Download Kinetic investigation of low-pH Fe(II) oxidation and development of a

Kinetic investigation of low-pH Fe(II) oxidation and
development of a method for Fe(III) regeneration as
part of a process aimed at H2S(g) removal
Research report submitted to Grand Water Research Institute
Ori Lahav and No'omi Levi
May 2009
Project No. 2008724
For the elimination of any doubt, it is hereby stressed that the staff member and/or the
Technion and/or the Technion Research and Development Foundation Ltd. Will not
be liable for any property damage and/or corporal damage and/or expense and/or loss
of any kind or sort that will be caused or may be caused to you or anyone acting on
your behalf, in consequence of this statement of opinion or this report, or in any
connection to it.
Copyright ©: Year of Publication: 2009; By: Ori Lahav and Noomi Levi and the
Technion Research and Development Foundation Ltd.
1
Table of contents
pp
Abstract
4
Abbreviations
7
Symbols
7
1. Introduction
8
1.1. Hydrogen Sulfide (H2S)
8
1.1.1. Properties of H2S(g)
8
1.1.2. Origins and concentrations of H2S(g)
8
1.1.3. Adverse effects of H2S(g)
10
1.1.4. H2S(g) threshold concentration
12
1.2. Methods for H2S(g) removal
13
1.2.1. Gaseous phase and solid phase processes
14
1.2.2. Physical-chemical processes in the aqueous phase
15
1.2.2.1. Regenerative gas scrubbing
18
1.2.2.2. LRSR process
18
1.2.3. Biological processes in the Aqueous phase
20
1.2.3.1. Biological methods
20
1.2.3.2. Microorganisms of particular interest for H2S(g) removal
21
1.2.4. The Fe(III)/Fe(II) LRSR process
22
1.2.4.1. Execution of LRSR processes at neutral to moderately
high pH levels
23
1.2.4.2. Execution of LRSR processes at low pH levels
1.3. The kinetics of Ferrous iron (Fe(II)) oxidation by oxygen
26
34
1.3.1. pH dependence
34
1.3.2. Fe(II) oxidation by O2 in natural waters
39
1.3.3. Fe(II) oxidation by O2 at low pH levels
41
1.4. Electrochemical oxidation of Fe(II)
44
1.4.1. Direct oxidation of Fe(II)
44
1.4.2. Indirect oxidation of Fe(II)
44
1.4.2.1. Competing reactions
45
1.4.2.2. Oxidation of Fe(II) by Cl2
46
2
Contents (continued)
1.4.3. Electrolytic production processes of active chlorine
46
1.4.3.1. Divided or undivided cell
46
1.4.3.2. Batch or flow-through mode
47
2. Hypothesis and objectives
50
2.1. Research hypothesis
50
2.2. Research objectives
51
3. Materials and Methods
52
3.1. Chemicals
52
3.2. Analytical equipment
53
3.3. Experimental
53
3.3.1. H2S(g) reactive-absorption experiments
53
3.3.2. Determination of practical precipitation potential of the working
solutions
54
3.3.3. Determination of Fe(II) oxidation rate
54
3.3.3.1. Catalytic oxidation of Fe(II)
54
3.3.3.2. Electrochemical oxidation of Fe(II)
55
3.4. Analytical methods
57
3.4.1. Determination of the total dissolved iron, phosphate and copper
concentrations
57
3.4.2. Determination of dissolved ferrous iron concentration
57
3.4.3. Determination of dissolved ferric iron concentration
57
3.4.4. Determination of chloride concentration
57
3.4.5. Determination of ammonia concentration
57
3.4.6. Analysis of species distribution by the MINEQL+ software
57
4. Results and Discussion
60
4.1. Catalytic oxidation
60
4.1.1. H2S(g) reactive-absorption efficiency
60
4.1.2. Practical precipitation potential in the working solution
62
4.1.3. Fe(II) oxidation rate in the catalytic oxidation experiments
64
4.1.3.1. Effect of pH
65
4.1.3.2. Effect of initial concentration of Fe(III)
68
3
Contents (continued)
4.1.3.3. Effect of initial concentration of Cu(II)
69
4.1.3.4. Effect of total phosphate concentration
73
4.1.3.5. Effect of total sulfate concentration
75
4.2. Electrochemical Fe(II) oxidation as part of the LRSR process
79
4.2.1. H2S(g) reactive-absorption efficiency
79
4.2.2. Practical precipitation potential in the working solution
83
4.2.3. Fe(II) oxidation rate in the electrochemical oxidation process
84
4.2.3.1. Direct electrooxidation
85
4.2.3.2. Indirect electrooxidation - effect of anode to cathode
surface area ratio (Sa:Sc)
85
4.2.3.3. Indirect electrooxidation - effect of current density
87
4.2.3.4. Indirect electrooxidation - effect of chloride concentration
89
4.2.3.5. Potential chlorine loss
91
4.2.3.6. Indirect electrooxidation - energy cost
91
4.2.3.7. Volume of solution in the electrolytic reactor
91
4.2.3.8. Electrode material
92
4.2.3.9. Changes in pH
92
5. Conclusions
93
Reference list
94
4
Abstract
Hydrogen sulfide (H2S) is a product of microbiological anaerobic activity. H2S is
emitted from environmental facilities and industrial processes such as petroleum
refineries, paper manufacturing, anaerobic digestion processes etc.
The removal of H2S from gaseous streams is required for the health of the general
public, occupational safety and for operational reasons.
The Liquid Redox Sulfur Recovery (LRSR) process is a common and promising
process for H2S(g) removal that is based on reactive absorption (i.e. absorption of the
gas into an aqueous solution followed by a chemical reaction). The H2S(g) is absorbed
into solution and then oxidized to S0 by an intermediate redox couple. The most
currently used couple is Fe(III)/Fe(II). Ferric (Fe(III)) oxidizes H2S(aq) to S0, and the
ferrous (Fe(II)) formed is oxidized back by oxygen to Fe(III). Working at pH7 to pH9
is favorable since the absorption of H2S(g) is efficient and the spontaneous oxidation of
Fe(II) by O2 is very rapid. However, the addition of organic chelates is essential in
order to avoid rapid precipitation of Fe(III) species as ferric oxides or oxy-hydroxides.
At low pH levels (pH~2) the solubility of ferric species is high enough to avoid the
need for chelating agents. Since the rate of Fe(II) oxidation is very low at this pH
range, bacteria populations have been suggested to catalyze the regeneration process,
mainly from the group Acidithiobacillus Ferrooxidans. However, this method suffers
from two major drawbacks: (1) the dependence of the process on the sensitive
autotrophic biomass and (2) the relatively rapid precipitation of Fe(III) solids, mostly
of the jarosite group.
The current work investigates two new approaches for applying the LRSR process at
pH1.0. The first method is based on catalytic oxidation of Fe(II) by O2 in the presence
of copper and phosphate, which are known to have a catalytic effect. The second is
based on electrochemical oxidation of Fe(II), directly (oxidation of Fe(II) on the
anode) or indirectly (oxidation of Fe(II) by chlorine formed on the anode). The very
low operational pH (pH1.0) was chosen in order to minimize precipitation of ferric
species for long-term LRSR operations. Two other factors to be considered in the
process are the efficient reactive-absorption of H2S(g) and rapid oxidation of Fe(II).
The catalytic oxidation method, although proven to accelerate Fe(II) kinetics in lab
tests, was found to be infeasible for the LRSR process, since despite of the low pH a
5
H2S introduced to the reactor precipitated with copper to form CuS, and Fe(III)phosphate solids also precipitated quite rapidly.
The direct electro-oxidation method was also found inapplicable since the efficiency
of the reactive-absorption of H2S(g) was very low in the absence of chlorides.
In contrast, the indirect electro-oxidation method was found to be highly feasible: the
reactive-absorption efficiency of H2S(g) was high in the presence of chlorides, there
was no appreciable amount of precipitates and the oxidation rate of Fe(II) was very
high.
6
Abbreviations
AC
Active chlorine
AD
Anaerobic digestion
A.F.
Acidithiobacillus Ferrooxidans
AFO
Animal feeding operations
DC
Direct current
DO
Dissolved oxygen
FAS
Ammonium iron(II) sulfate 6-hydrate
GRT Gas retention time
LRSR Liquid Redox Sulfur Recovery
ORP
Oxidation-reduction potential
Symbols
E
Efficiency of reactive-absorption
Fe(II)T
Total ferrous iron concentration
Fe(III)T
Total ferric iron concentration
Ox
The oxidized form of the redox
PT
Total phosphate concentration
R-
The reduced form of the redox
rFe ( II )
Rate of Fe(II) oxidation
rH 2 S( g )
Rate of H2S(aq) oxidation
ST
Total sulfate concentration
7
6. Introduction
The economic prosperity of any country is linked to the level of energy consumption.
Throughout the world, the increasing demand for energy is attributed to the growing
human population and subsequent desire to improve the living standard through
industrialization. The energy requirement for keeping pace with the development is
met through either conventional (coal, oil) or non-conventional (air, solar) energy
sources. Among those sources are gaseous fuels - combustible gases such as natural
gas (fossil origin) and biogas (anthropogenic origin). Gaseous fuels contain hydrogen
sulfide (H2S) in significant concentrations, along with other sulfur species, which
exist at traces concentrations. Different industries generate gaseous streams that
contain appreciable concentration of H2S (Pandey et al., 2004).
6.1. Hydrogen Sulfide (H2S)
6.1.1. Properties of H2S(g)
Hydrogen sulfide (H2S) is a flammable, colorless gas that smells like "rotten eggs" at
low concentrations. Sulfur containing compounds in general (such as mercaptants)
and H2S in particular are the products of the biodegradation of sulfur containing
amino acids such as cysteine and methionine by anaerobic bacteria. It is also a product
of anaerobic reduction of sulfate (SO42-) (Firer et al., 2008). H2S is soluble in various
liquids including water and alcohol (WHO, 2000). Henry’s constant for H2S is 0.102
M·atm-1 (Stumm and Morgan, 1996). In the aqueous phase H2S(g) is a diprotic weak
acid that dissociates according to the following equilibrium reactions (Noyola et al.,
2006):
H 2 S ( aq ) ↔ HS − + H +
HS − ↔ S 2 − + H +
pKa1 = 7.0
pKa 2 = 12.9
(1)
(2)
6.1.2. Origins and concentrations of H2S(g)
Only 10% of the total H2S global emission is of non-anthropogenic origin (Malhotra
et al., 2002), e.g. volcanic gases and hot springs (Busca and Pistarino, 2003). The
remaining is emitted from environmental facilities and industrial processes such as
petroleum refineries, coke production, viscose rayon production, the tanning industry,
wastewater treatment plants, paper and pulp manufacturing, food processing,
8
extraction of natural gas and anaerobic digestion (AD) processes (WHO, 2000;
Malhotra et al., 2002; Busca and Pistarino, 2003; Ramírez-Sáenza et al., 2009).
Sodium sulfide (Na2S) is one of the most widely used reagents in the tanning
Industry. Processing liquors possess sulfide ion concentrations up to 2000 mg/kg and
untreated tannery wastewater contain up to 20 mg/kg sulfide (Lawrence et al., 2000).
The content of sulfur in crude oils is typically in the range of 0.3–0.8 wt.%. A
systematic increase in the sulfur content of extracted crude oils has been recorded
over the past two decades and a further increase is anticipated (Lawrence et al., 2000).
For example, Pandey et al. (2004) recorded an H2S(g) concentration to be 52,600 ppm
in a refinery fuel gas sample. The composition of typical refinery fuel gas is presented
in Table 1.
Natural gas is a complex mixture containing desirable gaseous hydrocarbons and nonhydrocarbon components such as H2S, CO2 and water. According to Koros and
Mahajan (2000), the world market for natural gas is estimated at approximately 22
billion US$ annually. Around 40% of the approximately 54 trillion standard cubic feet
of gas reserves in the lower 48 United States are not being developed due to the high
cost of gas treatment (Koros and Mahajan, 2000). The H2S content of natural gas
ranges from 100 up to 300,000 ppm and is expected to increase (Lawrence et al.,
2000).
Anaerobic digestion processes are an effective technology for the reduction of the
organic matter and simultaneous production of energy through biogas production.
Biogas is a sub-product of anaerobic digestion that has a high energy value attributed
to its high methane content. The use of this gas as a fuel can decrease both energy
costs and operational costs in waste treatment plants where it is generated (Mesa et
al., 2004). It is also used in industrial-scale animal feeding operations (AFOs) to treat
animal wastes and attaining the same benefits. Biogas production and utilization is
constantly increasing, as it represents a “green”, renewable energy, obtainable in a
relatively economical way (Cosoli et al., 2008).
Biogas is generally composed of 60–65% methane (CH4) and 35–40% carbon dioxide
(CO2). Minor constituents include hydrogen sulfide (H2S), nitrogen gas (N2),
hydrogen gas (H2) and traces of oxygen (O2), carbon monoxide (CO), ammonia
9
(NH3), argon (Ar2) and other volatile organic compounds (VOC). The composition of
biogas depends on the type and concentration of organic matter to be digested, on the
physicochemical conditions in the digester (pH, alkalinity, temperature) and on the
presence of other electron acceptor species such as sulfates and nitrates (Noyola et al.,
2006). A typical composition of a biogas is given in Table 2.
Table 1. Characteristics of a refinery fuel
gas sample (Source: Pandey et al., 2004)
Component
Composition (vol.%)
Table 2. Typical composition of biogas
(Source: Pauss et al., 1987, quoted by ter
Maat et al., 2005)
Component
Composition (vol.%)
CH4
47.76
CH4
52–95
H2S
5.26
H2S
0.001–2
H2
1.61
H2
0.01–2
C2H6
17.87
CO2
9–45
Propane
11.12
N2
0.1–4
Propylene
4.88
O2
0.02–6.5
i-Butate
1.61
Ar
0.001
n-Butane
3.73
CO
0.001–2
Butylene
3.97
NH3
Trace
1.71
Organics
Trace
Pentane
4
1 vol.% = 10 ppm
As can be seen from Table 2 the range of H2S content in biogas vary considerably.
Concentrations in a similar range were reported also by Mesa et al. (2004), Chung et
al. (2006), Noyola et al. (2006), Qaisar et al. (2007), Cosoli et al. (2008) and
Fortunya et al. (2008). Much lower concentrations were measured in AFOs: Heber et
al. (2004, quoted by Gendel, 2007) measured H2S concentrations throughout the year
in swine facilities and obtained concentrations between 0.14 and 1.20 ppm. Okoli et
al. (2004) measured concentration of H2S in five intensive layer farms during the
month of August and obtained an average concentration of 1.53 ± 0.71 ppm.
10
6.1.3. Adverse effects of H2S(g)
As mentioned before, H2S has a very typical smell of rotten eggs and can be smelled
by the human nose at concentrations as low as 0.5 ppb (Busca and Pistarino, 2003;
Firer et al., 2008; Ramírez-Sáenza et al., 2009). It is the most characteristic bad odor
constituent in biogas and in the vicinity of anaerobic digesters and wastewater
treatment facilities in general (Noyola et al., 2006). At concentrations higher than 100
ppm, the olfactory system is affected and the human nose cannot sense the typical
smell (Busca and Pistarino, 2003; Firer et al., 2008).
Additionally, hydrogen sulfide is also highly toxic and may be lethal to mammals.
When inhaled, H2S(g) inhibits an enzyme which has an important role in mitochondrial
respiration (Firer et al., 2008). Exposure of human beings to low H2S concentrations
can cause headaches, nausea and irritation of eyes and throat as well as rhinitis,
Keratoconjuntivitis, photophobia, intense cough and bronchopneumonia. High
concentrations can cause paralysis of the breathing system, unconsciousness and
finally death (Merck, 1996 quoted by Noyola et al., 2006). Lethal doses, depending
upon exposure, can range from 300 to 1000 ppm. Clinical cases of sulfide poisoning
typically involve levels from 4 to 4,200 ppm. In Table 3 some health effects of H2S,
with respect to exposure concentrations, are presented. It should be mentioned that
H2S is highly toxic not only to mammals but also to aquatic species and vegetation.
Apart of being a nuisance and a health problem, the presence of H2S(g) in gaseous
fuels may cause operational difficulties. Combustion of a gaseous fuel that contains
H2S may lead to the formation of acid rain, due to oxidation of H2S(g) in combustion
to sulfur dioxide (SO2). Through heterogeneous oxidation on particulate matter, SO2
is further oxidized to SO3. The SO3 is highly soluble in water that results in the
formation of acid rain (Malhotra et al., 2002; Pandey et al., 2004).
H2S is a corrosive compound that attacks different materials - iron, copper, cement
etc. (Noyola et al., 2006). Thus, fuel gas with high content of H2S expedites corrosion
of engines, pipelines and biogas storage structures (Koros and Mahajan, 2000; Chung
et al., 2006; Fortunya et al., 2008)
Consequently H2S(g) has to be removed from gaseous fuels prior to combustion.
11
Table 3. Hydrogen Sulfide: established dose-effect relationships (Source: WHO,
2000)
H2S concentration
mg·m-3
ppm
1400 - 2800
1000 – 2000
effect
Immediate collapse with
paralysis of respiration
Reference
WHO, 2000a
Strong CNS stimulation,
750 – 1400
530 - 1000
hyperpnoea followed by
WHO, 2000a
respiratory arrest
450 – 750
320 – 530
210 – 350
150 – 250
70 – 140
50 – 100
15 - 30
10 - 20
Pulmonary oedema with risk of
WHO, 2000a; Firer
et al., 2008
death
Loss of olfactory sense
Serious eye damage
Threshold for eye irritation
WHO, 2000b
WHO, 2000b; Firer
et al., 2008
WHO, 2000b
a – adopted from: Hydrogen sulfide. Geneva, World Health Organization, 1981
(Environmental Health Criteria, No. 19).
b – adopted from: Savolainen, H. Nordic expert group for TLV evaluation. 40.
Hydrogen sulfide. Arbeta och hdlsa, 31: 1-27 (1982).
6.1.4. H2S(g) threshold concentration
The World Health Organization (WHO) recommends a H2S(g) guideline value of 2.1
ppm (0.15 mg/m3) for an average exposure time of 24 hours, to ensure no eye
irritation. Another recommendation of the WHO is a threshold of 10 ppb (7 µg/m3) for
a 30-minute average exposure period, to avoid odor annoyance among the exposed
populations (WHO, 2000).
The US Occupational Safety and Health Administration (OSHA) established an
acceptable ceiling concentration of 20 ppm at the workplace, with a maximum
concentration of 50 ppm for no more than 10 min. The US National Institute of
Occupational Safety and Health (NIOSH), and the Japanese and Swedish equivalents,
have set a maximum recommended exposure limit ceiling value (10 min) of 10 ppm
(Lawrence et al., 2000; Busca and Pistarino., 2003; Ramírez-Sáenza et al., 2009).
12
Considering the progressive nature of legislation and the pressure exerted by the
legislators as a result of public awareness, it is anticipated that increasingly stricter
regulations will be applied in the near future (Iliuta et al., 2004).
At present, no common standard has been defined for biogas upgrading to natural gas,
but it can be assumed that the European standards will be de-facto the guidelines for
most countries. Nevertheless, H2S concentration in biogas should be kept below 7
ppm (5 mg/m3) (Noyola et al., 2006). Normally the H2S specification for the product
gas will vary between 4 and 500 ppm H2S, depending upon further use (ter Maat et
al., 2005). A concentration of 4 ppm of H2S in natural gas must be reached before the
gas is put into a pipeline (DeBerry, 1997).
From all of the above it is clear that the removal of H2S from gaseous streams is
required for the health of the general public, occupation safety and operational
reasons. While concentrations as high as 50,000 ppm exist in gaseous fuels, threshold
concentrations are around 10 ppm.
6.2. Methods for H2S(g) removal
The removal of hydrogen sulfide from the gaseous streams depends on various factors
e.g., raw feed composition, treated gas quality, economic analysis of the process of
desulphurization, and the corrosion problems in existing operational units. The
gaseous streams, therefore, require desulphurization through a techno-economicallyviable process in order to meet the product purity requirement, to generate clean fuels
(in case of gaseous fuels) and also to conform to the stringent sulfur emission
standards (Pandey et al., 2003).
Most of the purification and treatment processes of gaseous streams that contain
H2S(g) result in generation of sulfate (SO42-) or elemental sulfur (S0) through
oxidation. Elemental sulfur is a non-corrosive solid that is easy to handle and
transport. In addition, it has a commercial value exceeding that of sulfuric acid,
although both are used in chemical processing and fertilizer production (Qaisar et al.,
2007). Most of the elemental sulfur is produced to-date in oil refineries and natural
13
gas treating plants, and is sold for the production of sulfur compounds such as sulfuric
acid (Busca and Pistarino, 2003).
Technologies for gaseous pollution control are varied. Fig. 1 shows the common
application of each type of technology based on pollutant concentration and air or gas
flow. Many of the H2S(g) removal processes are based on its absorption into an
aqueous solution followed by further treatment. Other processes are preformed in the
gaseous phase or involve a solid phase, as described in the following paragraph.
Figure 1. Applicability of various gaseous pollution control technologies based on
gas flow rates and concentrations to be treated (Source: Noyola et al., 2006).
6.2.1. Gaseous phase and solid phase processes
The most common gaseous phase process is the Claus process. It allows the recovery
of sulfur from highly concentrated streams, and is reported to be reliable at a sulfur
production rate higher than 20 tons per day (Busca and Pistarino, 2003). The process
is based on the oxidation of hydrogen sulfide to elemental sulfur, as described in Eq.
(3):
H 2 S + 12 O2 → S 0 + H 2O
(3)
The process consists of two steps (Eq. (4) and (5)):
step I: H 2 S + 32 O2 → SO2 + H 2 O
(4)
step II: 2 H 2 S + SO2 → 3S 0 + 2 H 2 O
(5)
14
The typical Claus process industrial configuration is comprised of two steps. In the
first step one-third of the H2S concentration is first oxidized to SO2 in a burner. A
second exothermic step is performed at lower temperatures, namely 473–573 K, in a
series of fixed bed catalytic reactors in the presence of alumina as a catalyst.
However, one step processes can also be performed. The produced sulfur is
condensed and stored (Busca and Pistarino, 2003). Numerous desulphurization plants
based on Claus and Super Claus processes are in operation worldwide (Pandey et al.,
2004). However, the need for high temperature makes this process economically
feasible mainly for systems that work already at high temperatures.
Other gaseous phase processes are thermal incineration and catalytic combustion,
which are used to remove highly concentrated VOC-containing streams, with a
benefit of additional energy supply. H2S(g) in the gas stream, if exist, will be oxidized
into sulfur dioxide (SO2), that has a much less pungent irritating odor with a much
higher recognition threshold (three orders of magnitude higher) than sulfide
compounds. Nevertheless, sulfur dioxide is also a toxic substance and its emissions
are regulated. The temperatures needed for these processes are high: above 1000K for
thermal incineration and 600-800K for catalytic combustion (Busca and Pistarino.,
2003).
Adsorption on solids is a common technology to recover volatile compounds from
contaminated air or gases. The pollutant molecules contact the surface of a solid
adsorbent and bond via weak intermolecular forces. The most used adsorbents for
H2S(g) removal are activated carbon and zeolite, as well as iron oxide based materials.
Adsorbent regeneration is accomplished by volatilization of the adsorbed compounds,
stripping (commonly with steam or nitrogen) or by purging with a solvent (Busca and
Pistarino, 2003; Noyola et al., 2006; Cosoli et al., 2008). The outlet stream from the
regeneration step contains much higher concentration of H2S than in the inlet stream,
which still needs to be annihilated.
6.2.2. Physical-chemical processes in the aqueous phase
Aqueous phase processes are usually preformed in near-ambient temperature, which
means that they require less energy than gas phase and solid phase processes.
15
Most of the H2S(g) removal methods in the aqueous phase are based on Reactive
absorption, i.e. absorption of the gas in an aqueous solution accompanied by a
chemical reaction. Absorption is usually done through scrubbing. In a scrubber,
transfer of pollutants from a gas stream to an aqueous phase is accomplished by
intense contact of the polluted gas with water or an absorbent solution, within a
packed column or a spray tower. Mass transfer depends on the concentration, the
air/water partitioning (Henry law) coefficient and the mass transfer resistance of the
scrubber system (Noyola et al., 2006). Scrubbing is reasonably applicable to gases
containing a high concentration of H2S to recover elemental sulfur through oxidation,
and is largely applied in refineries, for example, to recover H2S from gases arising
from hydrodesulphurization processes (Busca and Pistarino, 2003).
Different absorbing solution can be used, as presented hereunder.
Caustic scrubbing - Absorption is favored by highly alkaline conditions as can be
concluded from the dissociation constants of H2S(aq) (Eq. (1) and (2)). A gas stream
containing the pollutant is fed to an absorption tower with high alkalinity (i.e., NaOH
50% by weight, pH >12). The absorbent is not regenerated in this process, which
requires high reagent consumption and a proper final disposal of the spent solution
(Noyola et al., 2006).
Chemical precipitation with FeCl2 - The H2S contained in the gas stream is absorbed
in a scrubber with a solution of FeCl2 and the dissolved H2S is precipitated as FeS
according to the following reaction:
Fe 2+ + H 2 S → FeS ( s ) + 2 H +
(6)
The Fe+2 is not reused in this process, which means considerable reagent consumption
(Noyola et al., 2006).
Chlorine oxidation - After H2S has been absorbed in a scrubbing tower, it may be
oxidized with sodium hypochlorite to produce either elemental sulfur or sulfate,
depending on pH, according to the following reactions:
HS − + OCl − → S 0 + OH − + Cl −
−
HS + 4OCl − → SO 4
2−
at pH < 7.5
+ H + + 4Cl −
at pH > 7.5
(7)
(8)
Chlorine is not reused in the process, so it may result in a high operational cost.
Moreover, in the presence of organic compounds, chlorine oxidation is not an
16
attractive oxidation agent because of extended chlorine demand and the formation of
undesirable organic chloride compounds (Noyola et al., 2006).
Ozone oxidation - H2S or VOCs are dissolved in water within a scrubbing tower and
then they are oxidized by ozone. The H2S oxidation is practically instantaneous. Since
ozone is instable, its generation should be in situ, which may become an obstacle. In
addition, ozone is expensive (Noyola et al., 2006).
HS − + O3 → S 0 + OH − + O2
−
HS + 4O3 → SO4
2−
(9)
+ H + + 4O2
(10)
Potassium permanganate oxidation - After H2S is scrubbed, it can be oxidized using
potassium permanganate. This method is not attractive since it has a high cost and the
manganese oxide must be adequately disposed of to avoid a negative environmental
impact. Different sulfur compounds are produced depending on pH (Noyola et al.,
2006).
3H 2 S + 2 KMnO4 → 3S 0 + 2 H 2 O + 2 MnO2 + 2 KOH
at pH < 7.5
(11)
3H 2 S + 8KMnO4 → 3K 2 SO4 + 2 H 2 O + 8MnO2 + 2 KOH
at pH > 7.5
(12)
Hydrogen peroxide oxidation - The oxidation rate of sulfide with hydrogen peroxide
is relatively slow: 20 to 30 minutes contact time is normally required for a complete
reaction. The mechanisms of oxidation of H2S by hydrogen peroxide are not well
understood; however, it is suggested that direct oxidation of sulfide by hydrogen
peroxide depends on the reaction with oxygen released during gradual decomposition
of hydrogen peroxide (Noyola et al., 2006).
H 2 O2 + H 2 S → S 0 + 2 H 2 O at pH < 8.5
2 H 2 O2 + S 2− → SO2
2−
+ 2H 2 O
at pH > 8.5
(13)
(14)
All the removal processes described above may be highly effective, but all have a
common and significant drawback – a high reagent consumption, which usually
means high operational costs and the need for a proper disposal way for the waste.
Another promising group of techniques for hydrogen sulfide removal are the
processes that include regeneration of the reactive agent.
17
6.2.2.1. Regenerative gas scrubbing
In case of regenerative processes the reagent used to capture H2S can be recovered.
Different solvents that have a high affinity for H2S can be applied, such as ethanolamines seem (Busca and Pistarino, 2003; Fortunya et al., 2008). An example of this
process is described by the following reaction (Eq. (15)):
R 2 NH + H 2 S ↔ R2 NH 2 HS + heat
(15)
The heat produced by the exothermic reaction during the absorption step is used to
preheat the absorbent in the desorption step, since desorption is favored at higher
temperatures. By this means, the absorbent can be regenerated. However, H2S as well
as CO2 are discharged as exhaust product, so additional treatment processes would be
needed before final disposal (Noyola et al., 2006). Although these processes have
been extensively and successfully applied, they have many drawbacks such as high
energy and operating costs due to the regeneration of the absorbent phase (Fortunya et
al., 2008). Additionally, treatment with solvents like ethanol-amines leaves
concentration of residual H2S in the gas in the order of few ppm. This means that
deodorization is not achieved without further dilution or treatment (Busca and
Pistarino, 2003).
Another regenerative method is based on the precipitation of H2S(g) with metal ions
accompanied by subsequent regeneration of the solid metal sulfide formed, as
suggested by ter Maat et al. (2005). As mentioned earlier, the dissolved H2S molecule
acts as a diprotic acid. Sulfides of most bivalent metal ions, e.g. zinc, copper, silver,
lead, magnesia, nickel and tin are highly insoluble. Therefore aqueous solutions that
contain these metal ions can be used as a washing liquid in desulphurization
processes. The solid produced is separated from the spent absorbent, oxidized (into
CuO, for example) and then dissolved back in the spent absorption liquid.
Another, very common, regenerative scrubbing process is the LRSR process.
6.2.2.2. LRSR process
Liquid redox sulfur recovery (LRSR) processes offer inherently good hydrogen
sulfide removal and highly flexible operating properties at near-ambient temperatures
(nominally 20-50°C). The H2S is converted to elemental sulfur, that has a commercial
value, and the consumption of chemicals is minimal since the process is regenerative.
18
The overall reaction in LRSR processes is the same as for Claus process H2S
conversion (Eq. (3)) (DeBerry, 1997). However, direct oxidation of H2S by oxygen at
ambient temperatures is slow and side reactions tend to form undesirable sulfur
oxyanions as byproducts. Therefore, oxidation of H2S in the LRSR absorber is carried
out by an intermediate redox couple (“catalyst”) denoted Ox (for the oxidized form of
the redox couple) as described in Eq. (16).
2Ox + H 2 S ( aq ) → S 0 + 2 R − + 2 H +
(16)
The symbol R- is used for the reduced form of the redox couple, which is a one
electron redox couple in this example. The reduced form of the redox couple
produced in the reaction above is regenerated with air in the oxidizer, according to the
overall reaction:
4 R − + O2 ( g ) + 4 H + → 4Ox + 2 H 2 O
(17)
As shown, this reaction also consumes hydrogen ions, neutralizing the hydrogen ions
generated in the sorption step (Eq. (16)). The sulfur slurry is removed from the
oxidizer and sent to a filtration unit for removal as a wet (approximately 50% solids)
cake (DeBerry, 1997).
As the catalyst for the LRSR process the following redox couples can be used:
V5+/V4+; Fe3+/Fe2+; Co3+/Co2+ and As5+/As3+. Due to high oxidation rates of H2S by
Fe3+ (DeBerry, 1997; Hua et al., 2001) and the fact that the cations of vanadium (V),
cobalt (Co) and arsenic (As) are toxic and dangerous to the environment, the iron
couple appears to be the most suitable for the LRSR process. The use of vanadium
(the so-called Stretford process) has dominated in the 1980's, but stopped spreading
because of the chemical degradation and environmental problems associated with
vanadium. The iron based processes is gradually replacing the vanadium systems
(Hua et al., 2001). More than fifty plants based on Vanadium catalyst are in operation
for desulphurization of gaseous streams containing H2S. Similarly, around forty plants
all over the world based on iron based liquid-redox processes are in operation (Pandey
et al., 2004).
The continuing search for more economical methods has led to investigations into
microbiological solutions for purifying H2S-containing gases, as well as for the
desulfurization of coal and petroleum. Biological methods for H2S(g) removal are
discussed in the next section.
19
6.2.3. Biological processes in the Aqueous phase
Biological processes are characterized by low capital costs and low energy
requirements (Malhotra et al., 2002; Pandey et al., 2003; Ramírez-Sáenza et al.,
2009), since they can proceed around ambient temperatures and at atmospheric
pressure (Qaisar et al., 2007). A variety of biochemical processes using various
bacterial species are reported in the literature that are capable of removing sulfur in
various forms from gaseous emissions. In these processes, the microorganisms
metabolize the H2S, as a source of energy for growth and maintenance, producing
sulfate or elemental sulfur (Malhotra et al., 2002; Noyola et al., 2006).
In order to make a good selection of a treatment method, flow rate, type of pollutant
and its concentration must be considered. Bio-filters, bio-scrubbers and bio-trickling
filters have been proven to be a suitable, environmentally friendly and cost-effective
alternative for waste gas treatment, especially for the treatment of low concentrations
of H2S (Fortunya et al., 2008).
6.2.3.1. Biological methods
Biofilters – The contaminated gas is continuously fed into the biofilter, while a
nutrient solution is discontinuously added. The filter bed contains microorganisms
growing as a biofilm on the surface and crevices of the support. The H2S(g) is
absorbed from the gas to the aqueous bio-film and subsequently oxidized by
appropriate bacteria populations. The support media should have high porosity, high
buffer capacity, high nutrient availability, and especially high moisture retention
capacity (Qaisar et al., 2007). This technique is becoming widely accepted due to the
high processing efficiency at low sulfur concentrations, the moderate capital costs,
and the very low maintenance costs. Mixed micro-organism cultures naturally grow
on appropriate natural biofilter beds so abatement of many volatile compounds can be
obtained simultaneously (Busca and Pistarino., 2003).
Biotrickling filters - The operational principle of a biotrickling filter is similar to that
of a biofilter. In this device, polluted air is passed through a packed non-submerged
column where liquid is continuously down-flow recirculated through the packing. The
pollutant is first absorbed in the falling liquid film and transferred to the
microorganisms that grow attached to the surface of these supports. The liquid
provides moisture, nutrients, pH control to the biofilm and allows the removal of
20
inhibiting products. Eventually, excess biomass is sloughed off by the trickling liquid
and stable operation can be achieved (Noyola et al., 2006).
Bioscrubbers - Firstly H2S(g) from the gas phase is absorbed in a recycling water
stream followed by biological oxidation of H2S in the liquid. Nutrient addition and pH
are continually controlled in the bioreactor in order to maintain microbial growth and
high activity. The excess biomass and byproducts are purged from the system (Noyola
et al., 2006; Qaisar et al., 2007).
Membrane bioreactors - In a membrane bioreactor, the H2S(g) is transferred through a
membrane to a biofilm attached on the other side of the barrier, where nutrients and
oxygen are provided. A distinct characteristic of membrane bioreactors is the fact that
the polluted gaseous stream and the biomass is physically segregated which allows the
use of waste gas treatment in certain extreme applications such as indoor air (Noyola
et al., 2006).
6.2.3.2. Microorganisms of particular interest for H2S(g) removal
The use of microorganisms to oxidize H2S, producing sulfate or elemental sulfur as a
consequence of complete or incomplete metabolism, respectively, has been
considered a potential alternative for application on a large scale (Zhang and Tong,
2006). Among the H2S oxidizing microorganisms, Thiobacillus seems to be
particularly suited for engineering applications due to its simple nutritious
requirements, its high effectiveness and resistance to toxic substances and the wide
pH range it can tolerate (Noyola et al., 2006).
The most common reaction is a direct oxidation of sulfide to sulfur and sulfates by
means of oxygen provided by air (obligate autotrophs). In other cases (Thiobacillus
denitrificans) nitrate reduction to N2 allows the oxidation of sulfide to sulfate.
Particularly, Acidithiobaillus ferroxidans raises a very simple and effective process
for H2S treatment in which the oxidant is regenerated by the microorganisms (as part
of the Fe3+/Fe2+ LRSR process) (Noyola et al., 2006). Certain photosynthetic bacteria
belonging to families Chromatiaceae and Chlorobiaceae are also being used to
metabolize H2S effectively (Malhotra et al., 2002; Mesa et al., 2004).
The conventional physicochemical methods for removing hydrogen sulfide from
gaseous streams require large investment and operational costs e.g. high pressures,
high temperatures or special chemicals. Due to this high costs pre-treatment of
21
gaseous fuels contributes significantly to the overall operation and maintenance costs
of any energy recovery system. In addition, secondary hazardous wastes are generated
in most processes (poisoned catalysts, contaminated reactor liquids and corroded
reaction vessels), which need to be treated. (Malhotra et al., 2002; Pandey et al.,
2004; Qaisar et al., 2007). Microbiological processes, on the other hand, tend to be
efficient only at low H2S(g) concentrations. Additionally, biological systems need to be
fed continuously with nutrients and humidity (Mesa et al., 2004; Chung et al., 2006).
If the wanted microbiological population is washed out for some reason (e.g., blowdown of the system, strengthening of another population) it takes time to redevelop
the population.
From all the H2S(g) removal methods reviewed above, the Fe2+/Fe3+ LRSR process
seems to be more cost effective than the other physical-chemical processes and more
reliable than the common biological processes. The current research focuses on the
Fe3+/Fe2+ LRSR process as an efficient method for H2S(g) removal.
Note that from this point onward the total-iron species in the oxidation state of 3+ and
2+ will be addressed as ferric (Fe(III)) and ferrous (Fe(II)) respectively. The signs
Fe3+ and Fe2+ will represents the non complexed ions only.
6.2.4. The Fe(III)/Fe(II) LRSR process
As mentioned earlier, the iron couple is advantageous for the LRSR process because
of its high reaction rate and low toxicity (compared with vanadium, cobalt and arsenic
couples).
The H2S(aq) is oxidized by Fe(III) to elemental sulfur according to the following
equation:
2 Fe( III ) + H 2 S ( aq ) → S ( s ) + 2 Fe( II ) + 2 H +
(18)
The regeneration of Fe(III) is basically done by oxidation of the Fe(II) by O2:
2 Fe( II ) + 12 O2 + 2 H + → 2 Fe( III ) + H 2 O
(19)
In order to achieve an efficient and continuous process, the flux of ferrous should be
similar to the flux of ferric. The flux rates depend on the generation rates of Fe(II) and
Fe(III) in the process. Solution pH is often a key variable in LRSR processes, since it
has a major effect on the generation rates of Fe(II) and Fe(III) and consequently on
other parameters of the process.
22
6.2.4.1. Execution of LRSR processes at neutral to moderately high pH levels
Hydrogen sulfide is the acidic species of the sulfide weak-acid divalent system (Eq.
(1) and (2)), thus its absorption is preferable in alkaline solutions. In addition,
spontaneous oxidation of Fe(II) is very slow below pH 6 (Stumm and Morgan, 1996).
The downside is that at pH levels above pH3 Fe(III) tends to rapidly precipitate as
iron oxides or iron oxy-hydroxides. The use of organic chelating agents, which have
the capacity to bond with cationic ions, proved to be efficient to prevent iron
precipitation over a wide pH range, typically up to 10, and to afford satisfying
regeneration rate of Fe(III) (Iliuta et al., 2004). The main chelating agents in
commercial use are amino and polyamino-polyacetic acids e.g. nitrilo-triacetic acid
(NTA), ethylene dinitrilo-tetracetic acid (EDTA) and N-(2-hydoxyethyl) ethylenediamine-N,N',N' triacetic acid (HEDTA) (Hua et al., 2001).
DeBerry (1997) concluded that iron-chelate processes are probably best operated
above pH7 from the standpoint of the initial Fe-S complex formation and sulfur
formation. Other factors to be considered in an LRSR process with chelating agents
are detailed below.
H2S reactive-absorption efficiency - For good absorption efficiency and small
scrubber size the rate of reaction of the redox catalyst with H2S should be as high as
possible. The more successful catalysts appear to rapidly form a complex with H2S
and then more slowly convert it to sulfur. The initial complex formation reduces the
equilibrium backpressure of H2S. The initial reaction rates of H2S with several
oxidizing agents are shown in Table 4.
Table 4. Rates of some oxidizing agents with Hydrogen sulfide (Source: DeBerry,
1997)
Oxidant
O2
H2O2
Second order rate constant k (M-1sec-1) Comments / source
1.5 × 10-3
For HS-+O2HS+O2-;
20ºC; Resch et al., 1989
-1
Resch et al., 1989
-1
4.5 × 10
V(V)
4.2 × 10
20ºC; Radian
FeOH2+
1.5 × 103
20ºC; pH2; Asai et al., 1990
Fe(III)EDTA
6.2 × 103
20ºC; Radian
23
Also important is the capacity of the scrubbing liquid for H2S. Use of a concentrated
redox catalyst can increase the capacity of the sorbent for H2S. However, high catalyst
concentrations could also lead to fast precipitation of sulfur in the scrubber and
consequent plugging (DeBerry, 1997). Control of sulfur formation is a key topic, as
discussed hereunder.
Sulfur Formation - The formation of solid sulfur is a key step in LRSR processes. The
efficient formation of a separable solid product is important to economical continuous
operation of these processes. However, premature formation of sulfur in the H2S
scrubber can cause plugging of the scrubber internals (internals are necessary to get
good gas/liquid mass transfer and high removal efficiency). Faster oxidation of H2S
has been reported to lead to faster precipitation of S0, and hence to smaller particle
size of the S0 precipitates and settling difficulties (DeBerry, 1997).
Redox Catalyst Regeneration - The reduced form of the redox catalyst must be
converted back to the oxidized form in the regenerator, which is usually an airsparged vessel. This vessel is usually large compared to the other vessels used in
LRSR processes. In addition, the air compressors and energy needed to run the
compressors can be significant. Thus, rapid regeneration reaction rates are needed for
economical operation. Generally speaking, the higher oxidation potential the
oxidizing catalyst has, the slower and more difficult will be the regeneration step. This
is because as the redox potentials get more positive, they approach the redox potential
of oxygen, resulting in a decreased driving force for the regeneration of the redox
catalyst by O2. Thus, the catalyst redox potential should be high enough to give
efficient H2S scrubbing (as discussed above) but low enough to allow efficient
regeneration by air. The chelate molecule plays a major role in the rate of regeneration
of Fe(III) from Fe(II) in the chelated iron LRSR processes. The air oxidation rate of
uncomplexed ferrous ion (Fe2+) is many orders of magnitude slower than that of Fe2+
chelated with EDTA or similar chelating agents. However, Fe(III) chelated with
EDTA is a much weaker oxidizing agent than simple aquated (complexed only by
water molecules) Fe(III). Oxidation rates of several compounds by oxygen are
compared to the rate of oxidation of Fe(II)EDTA in Table 5. The second order rate
constants were derived primarily from the overall or the slowest step.
The rate constants in Table 5 range over six orders of magnitude. Note that both the
uncatalyzed oxidation rate of H2S by O2 and the rate of uncomplexed Fe(II) oxidation
by O2 with are very low (DeBerry, 1997).
24
Table 5. Comparison of rates of oxidation of species by Oxygen (Source: DeBerry,
1997)
Reactant
H2 S
Second order rate constant k [M-1sec-1] Comments / Adopted from
1.5 × 10-3
Resch, 1989
pH 8.8; Radian*
V(IV)
1.0
Fe2+(aq)
4.0 × 10-4
Lacey, 1970
Fe(II)EDTA
2
2.8 × 10
Radian*
Fe(II)DTPA
1.4 × 101
Radian*
Fe(II)HEDTA
6.0 × 102
Radian*
*
Radian Corporation, Austin, Texas
Consumption of Chemicals - There is several sources of chemical consumption in
liquid redox processes. One of these, carry-out of chemical reagents with the sulfur
product, is in large part a “physical” process, although some chemical binding to the
surface of the sulfur may occur. A major problem for iron chelate-based processes is
chemical degradation of the chelating agent. The degradation is generally thought to
be caused by free OH● and H2O2 radicals or other highly reactive intermediates
generated in the air oxidation reactions. A number of compounds have been tested in
an attempt to scavenge the hydroxyl radicals and thus protect the chelating agents.
Thiosulfate (S2O32-) is known as an effective scavenger for the hydroxyl radicals and
has been proposed as a prospective degradation inhibitor. However, most of these
scavengers are consumed more or less rapidly and thereby lead to a buildup of soluble
salts in solution. This will eventually necessitate a blow-down or a solution cleanup
step. Eventually, the addition and consumption of chelate agents and radical
scavengers elevate the cost of the LRSR process (DeBerry, 1997; Hua et al., 2001).
Another source of loss is blow-down of solution necessitated by byproduct formation
(see below).
Byproduct Formation Rate - Oxidation of H2S beyond the oxidation state of elemental
sulfur results in the formation of soluble sulfur oxyanions which are difficult to purge
from the LRSR process stream. If these salts reach high concentrations, they can
cause scale formation and operating problems. The salts must be removed by “blowdown” or chemical treatment of the solution which results in loss of some of the
catalyst and economic penalties. One of the main sources of this problem is the slow
25
and/or incomplete oxidation of sulfide, which leads to elevated concentrations of
polysulfides in solution. These polysulfides are readily converted to sulfur oxyanions
when the solution contacts air in the regenerator. Chelated iron processes, with respect
to vanadium/ADA processes, have much faster sulfide-to-sulfur conversion kinetics,
and the concentration of polysulfide in these solutions is usually quite low.
Paradoxically, one of the soluble sulfur oxyanion byproducts, thiosulfate, is a
desirable constituent of current liquid redox processes since it is a very good inhibitor
of chelate degradation, as described above. Thus the control and management of
thiosulfate formation is a part of the operating protocol of such processes (DeBerry,
1997).
To summarize, the use of organic chelating agents allows faster kinetics for
conversion of H2S to sulfur and faster regeneration of the catalyst with oxygen with
respect to earlier liquid redox processes. One unfortunate byproduct of this
evolutionary step is catalyst degradation, which is controlled in large part by
inhibitors and management of solution chemistry. However, most of these inhibitors
are consumed rapidly. Another byproduct of the faster kinetics associated with the
new catalysts is difficulty in controlling the formation of solid sulfur.
In order to avoid the use of chelates and radical scavengers it is possible to perform
the LRSR process with the iron couple at low pH levels. This is discussed in the next
paragraph.
6.2.4.2. Execution of LRSR processes at low pH levels
At low pH levels (pH1 to pH2) the solubility of ferric species is high enough to avoid
the need for chelating agents.
Since the rate of ferrous oxidation is very low at this pH levels (see Subsection
1.2.4.1), bacteria are added to catalyze the regeneration process. The most commonly
used bacteria are Acidithiobacillus Ferrooxidans (A.F.). These are a gram-negative
rod-shaped, chemo-autotrophic, mesophilic and acidophilic bacteria, with the ability
to oxidize Fe(II) in acidic solutions and couple the energy to support carbon dioxide
fixation and growth. This ability is particularly suited for regeneration of Fe(III) in the
LRSR process. The A.F. bacteria grow at pH values between pH 1 and 4.5, and the
optimal pH and temperature ranges for ferrous iron oxidation are 1.8–2.5 and 30 –35
°C, respectively (Malhotra et al., 2002; Mesa et al., 2004; Chung et al., 2006). In
26
Addition, the problems associated with other H2S oxidation microbiological processes
are avoided as H2S does not have an inhibiting effect on A.F. and SO42- is not
accumulated in the medium (Noyola et al., 2006 ; Ramírez-Sáenza et al., 2009).
The biological LRSR process at low pH levels consists of two steps, usually done in
two different reactors. The first one is a scrubbing reactor, where H2S(g) is absorbed
into ferric sulfate (Fe2(SO4)3) solution and oxidized to S0, while Fe(III) is reduced to
Fe(II) (Eq. (20)). The solution from the first reactor flows to the second reactor, an
aerobic bioreactor, where Fe(III) is regenerated from Fe(II) by A.F. bacteria (Eq.
(21)). The ferric iron is then recycled into the first reactor to repeat the cycle
(Malhotra et al., 2002; Ebrahimi et al., 2003; Mesa et al., 2004; Pandey et al., 2004;
Chung et al., 2006).
H 2 S + Fe2 ( SO4 ) 3 → S 0 + 2 FeSO4 + H 2 SO4
(20)
2 FeSO4 + H 2 SO4 + 0.5O2 → Fe2 ( SO4 ) 3 + H 2 O
(21)
The bioreactor can be operated either as a fixed or fluidized bed. The process is
schematically depicted in Fig. 2.
Figure 2. Process scheme of a chemo-biological process for H2S(g) removal (Source:
Ebrahimi et al., 2003)
A distinct advantage of the process is that the reaction of H2S with Fe2(SO4)3 is so
rapid and complete that it does not produce toxic waste or generate other waste
products apart from sulfur (Ebrahimi et al., 2003; Mesa et al., 2004). The overall
process rate is of a similar order of magnitude as that of the alkaline LRSR process
with chelating agents (Hua et al., 2001). Hence, the operating cost of the chemobiochemical processes is extremely low as compared with other chemical redox
processes for desulphurization of the gaseous fuels (Pandey et al., 2003). In general,
27
the operation costs of this process are around one third of those of conventional
processes such as scrubbers and absorption columns (Noyola et al., 2006).
In order to achieve an efficient and continuous process, the flux of Fe(II) from the
absorbent reactor should be similar to the flux of Fe(III) from the bioreactor.
The Fe(II) generation step in the absorption reactor depends on two reactions
occurring one after the other: H2S(g) absorption and H2S(aq) oxidation (i.e., reactive
absorption). Thus the Fe(II) generation is related to the efficiency of both reactions.
Consequently, the main factors that affect Fe(II) generation, as reviewed in literature
are: the H2S(g) concentration, the Fe(III) concentration and pH. These factors are all
reported to have a positive effect on the Fe(II) generation rate.
The positive effect of H2S(g) flux on Fe(II) generation was reported by Asai et al.
(1990), Ebrahimi et al.(2003), Pandey et al. (2003), Chung et al. (2006) and Gendel
(2007). Most of these works found a linear relationship between H2S flux and ferrous
generation rate, as shown in Fig. 3.
Figure 3. Relationship between inlet H2S(g) loading and Fe(II) production rate
(Source: Chung et al., 2006).
The positive effect of Fe(III) was reported by Asai et al. (1990), Ebrahimi et
al.(2003), Pandey et al. (2003), Mesa et al. (2004), Pandey et al. (2004), Chung et al.
(2006) and Gendel (2007). Asai et al. (1990) reported that the species which reacts
with H2S is FeOH2+ and that the reaction is irreversible and first order in both H2S and
FeOH2+. The overall reaction is:
H 2 S + 2 FeOH 2 + → S 0 + 2 Fe 2 + + 2 H 2 O
28
(22)
which consist of the following steps:
step I: H 2 S + FeOH 2+ → H 2 S • FeOH 2+
(23)
step II: H 2 S • FeOH 2+ + FeOH 2+ → S 0 + 2 Fe 2+ + 2 H 2 O
(24)
Asai et al. (1990) assumed that FeOH2+ was the sole reactive species in the system.
However, Ebrahimi et al. (2003) have evaluated the concentration of all possible
ferric species by taking into account all of the ferric complex equilibrium equations
and showed that other ferric hydroxide species (Fe3(OH)45+ and Fe2(OH)24+) were on
the same order of magnitude as FeOH2+ concentration under the experimental
conditions (Fig. 4). Thus Ebrahimi et al. (2003) concluded that the reactive-absorption
rate of H2S(g) was first order with respect to both H2S(g) and the total ferric iron
concentration.
Figure 4. Concentrations of different ferric species as a function of pH at 45ºC and
0.2 M Fe2(SO4)3 (Source: Ebrahimi et al., 2003)
Pandey et al. (2003) found that further increase in the H2S(g) loading rate resulted in
the decrease of its removal efficiency (and consequentially Fe(II) generation) due to
limited availability of Fe(III) in the reaction medium. It is also possible that at high
29
Fe(III) concentration the H2S(g) concentration becomes the limiting factor. This may
explain the decrease in Fe(II) generation with further increase in Fe(III) concentration,
as reported by Ebrahimi et al.(2003) , Pandey et al. (2003), Mesa et al. (2004) and
Pandey et al. 2004). Another explanation is that above a certain concentration of
Fe2(SO4)3 the ionic strength and viscosity of the solution increase remarkably
resulting in a decrease of the solubility, diffusivity and liquid side mass transfer
coefficient of H2S and hence absorption rate (Ebrahimi et al., 2003). In the
experiments conducted by Ebrahimi et al. (2003) the limiting concentration was 0.3
M. However, this value is different in each work, due to different conditions of
experiments.
All experiments reviewed were conducted at pH between 1.1 and 2.3. A positive pH
effect was reported by Asai et al. (1990), Ebrahimi et al.(2003), Pandey et al. (2003)
and Pandey et al. (2004). There are two possible causes: first, with the increases in pH
there is an increase in ferric hydroxide species that have a positive effect on Fe(II)
generation according to Ebrahimi et al.(2003); second is the poor solubility of H2S in
low pH ferric sulfate solutions (Pandey et al., 2003).
Two other parameters that affect H2S(g) reactive-absorption and can be modified quite
easily are the surface area for gas-liquid mass transfer and the gas retention time in the
adsorbing solution. Pandey et al. (2003) investigated the reactive-absorption
efficiency of H2S(g) in a packed bed reactor and found efficiencies above 96%.
Packing material that had large surface area resulted in high removal efficiency of
H2S(g) due to better mass transfer of H2S from the gaseous bulk to the liquid phase.
However, the sulfur produced in the process with the packing material with the largest
surface area accumulated in the chemical oxidation unit, which resulted in plugging of
the reactor. Chung et al. (2006) examined three gas retention times and their results
indicated that a longer GRT could elevate H2S(g) removal efficiency. Gendel (2007)
reached similar results when he examined two different flow rates of the treated air,
and found that at low flow rate the removal efficiency of H2S(g) was higher.
As for the Fe(II) generation step in the bioreactor - in all the works that were reviewed
an efficiency higher than 90% for Fe(II) oxidation by A.F. was achieved (Malhotra et
al., 2002; Pandey et al., 2003; Mesa et al., 2004; Pandey et al., 2004; Chung et al.,
2006 and Gendel, 2007). An Fe(II) oxidation rate of at least 11.3 mM/h (0.63 g/l/h)
was reported by Mesa et al. ( 2004), Chung et al. (2006) and Gendel (2007).
30
A.F. can convert ferrous ions to ferric ions in the presence of TDS concentration up to
50,000 mg/l and even 60,000 mg/l with proper acclimatization (Malhotra et al., 2002).
The presence of elemental sulfur hardly affects the bio-oxidation of Fe(II) to Fe(III)
according to Malhotra et al. (2002). When the dissolved oxygen concentration is not a
growth limiting substrate i.e., above 0.29 mg/l (Liu et al., 1988, quoted by Gendel,
2007) and for a given pH and temperature, both A.F growth and the ferrous iron
oxidation rate depend on three parameters: substrate (i.e. ferrous iron) concentration,
product (i.e. ferric iron) concentration (competitive and non-competitive inhibitions)
and the biomass concentration.
The ratio between the volumes of the chemical (absorption) and biological reactors
depends on the ratio between the rates of H2S(g) removal and Fe(II) oxidation.
According to Gendel (2007), who examined the removal efficiency of H2S(g) in
concentrations between 2 and 57 ppm, the required volume for the biological reactor
can be designed three orders of magnitudes smaller than the chemical absorption
reactor. However Chung et al. (2006), who experimented with H2S concentrations of
500-1500 ppm, conclude that the volume ratio of biological reactor to chemical
reactor should be 13.5:1.
All in all, the method of biological regeneration of Fe(III) by Acidithiobacillus
Ferrooxidans as part of the LRSR process at low pH suffers from two major
drawbacks: (1) the dependence of the process on the sensitive autotrophic biomass
and (2) the relatively rapid precipitation of Fe(III) solids at the pH range optimal for
A.F. bacteria (mostly as ferric sulfate-hydroxides of the jarosite group).
With respect to the first point, it was mentioned earlier (Subsection 1.2.3.2) that
biological systems need to be fed continuously and it takes time to redevelop a
bacterial population. This is true especially for autotrophic bacteria, as A.F., which are
much more sensitive to changes in environmental conditions (e.g. temperature,
fluctuation in substrate flux, pH) and have a very low yield coefficient. To exemplify
the time needed for the development of A.F. population: in the experiment conducted
by Chung et al. (2006) the A.F. bacteria that were used for the experiment were
isolated from acid mine drainage, were grown in a growth medium for 7 days and
then were cultivated for two further months on the packing material (granular
activated carbon).
31
As for the second point, the formation of ferric iron precipitates has been reported in
most studies that involve ferrous iron oxidation by A.F. In many of the reports the
precipitates were identified as the iron hydroxide jarosite, a basic ferric hydroxylsulfate with the general formula MFe3(SO4)2(OH)6, where M may stand for K+, Na+,
NH4+ or H3O+ (Pagella and Faveri, 2000; Mesa et al., 2004). The formation of
jarosites depends strongly on pH. Gendel (2007) showed that the mass of precipitate
that was formed increased with an increase in the initial pH from 1.25 to 2.5.
Although the formation of jarosite was significantly reduced at lower pH values,
jarosite precipitation did not cease and remained a major problem for long-term
operation even at low pH values. A further decrease in pH (i.e. pH≤1.0), may
diminish the problem, however at such pH the oxidation efficiency of Fe(II) by A.F. is
very low, and the bacteria require a long acclimation period.
The problems associated with jarosite precipitation are the following: (1) The
precipitation reduces the soluble ferric concentration in the aqueous phase, and thus
less ferric is available for oxidizing H2S in the desulphurization process; (2)
Precipitation of ferric compounds tends to block pumps, valves, pipes etc., and causes
maintenance difficulties; (3) Precipitation of ferric compunds on the surface area of
the carrier in fixed biological beds reduces both the amount of surface area available
for biomass growth, and the flux of the substrate (ferrous iron) into the biofilm, thus
reducing the overall efficiency of the biological step. In experiment by Mesa et al.
(2004) the final precipitate was found to accumulate on the support at rates ranging
between 0.1 and 0.8 g/l/h. This implies that a one cubic meter reactor working
continuously for 100 days would accumulate approximately 2000 kg of precipitate.
Jarosite formation depends also on the ferric iron concentration: when the
concentration is below 1000 mg/l jarosite precipitation is considerably slow (Gendel,
2007). However, at low ferric concentration the reactive-absorption efficiency of
H2S(g) is low.
Only one study was found to describe an LRSR process at low pH, without the use of
Acidithiobacillus Ferrooxidans for Fe(III) regeneration. Zhang and Tong (2006) used
copper to catalyze Fe(II) oxidation by O2. The experimental conditions were as
follows: 0.54 M Fe(II), 1.07 M Fe(III), 1.26 M Cu(II); pH was low but not specified;
H2S(g) concentration was around 1000 ppm and the gas (air + H2S) flow was 11.7
32
m3/min (3.0 l/m3 liquid-gas ratio). According to Zhang and Tong (2006) the cupric ion
precipitated rapidly as cupric sulfide (Cu(II)S(-II)) (Eq. (25)), however it was then
rapidly oxidized by Fe(III) to elemental sulfur (Eq. (26)). The Cu2+ ion has a catalytic
effect on Fe(II) oxidation, hence the oxidation of Fe(II) and consequently the
regeneration of the absorbing solution were fast.
(25)
Cu 2+ + S 2− → CuS
(26)
2 Fe 3+ + CuS → S 0 + 2 Fe 2+ + Cu 2+ quick
quick
Zhang and Tong (2006) reported that the H2S(g) removal efficiency of their suggested
process was almost 100% during the entire experiment (that lasted 60 minutes) while
the concentrations of cupric ion and ferric ion were very close to the initial values. No
CuS was found to accumulate. In addition, sulfur selectivity after 4 h of experiment
was calculated to be higher than 96% (Sulfur selectivity = Moles of sulfur produced
per moles of H2S reacted). This process may be a successful alternative for an LRSR
process at low pH, without the disadvantages of a biological system.
To sum up all of the above, the methods for H2S(g) removal are varied, and the favored
method depends on the gas origin and content. The aqueous phase methods usually
have the lowest energy requirements, and among them the Fe(III)/Fe(II)-LRSR
method seems favorable since the oxidizing agent is regenerated in the system and the
only products are H2O and S0. When conducted at neutral to high pH (7-9) the H2S(g)
removal is very fast, but the use of chelating agents is essential in order to minimize
precipitation, mainly of ferric hydroxides. These chelating agents tend to degrade with
time, a fact which eventually leads to high operation costs. It is possible to perform
the Fe(III)/Fe(II)-LRSR method at low pH levels, in order to minimize precipitation,
with the use of Fe(II) oxidizing bacteria. In this way the use of chelating agents is
avoided, but the process becomes dependent on a very sensitive biomass (autotrophic
bacteria). Moreover, some precipitation still occurs, which is critical in continues
processes.
In this work two alternatives were investigated for Fe(III) regeneration at low pH
levels. The first one is chemical oxidation of Fe(II) by O2 in the presence of catalytic
reagents and the second is the use of electrolysis to oxidize Fe(II). In both methods
the pH would be kept around pH1.0. Thus, the two major drawbacks of the low pH
biological LRSR are bypassed: there is no dependence on a biological system and the
33
precipitation problem becomes minor. Before describing the methods in more details
the mechanisms of ferrous oxidation by O2 and electrochemical oxidation are
extensively reviewed.
6.3. The kinetics of Ferrous iron (Fe(II)) oxidation by oxygen
6.3.1. pH dependence
As early as 1906, the rate of oxidation of ferrous ion by dissolved oxygen in near
neutral pH solutions was shown to be (Goto et al., 1970):
−
d [ Fe( II )]
= k[ Fe( II )][OH − ] 2 PO2
dt
where k = 8.0(± 2.5) × 1013 min −1 atm −1 mol −1liter 2
(27)
at 20 0 C .
This kinetic equation was confirmed empirically and used in many studies (Sung and
Morgan, 1980; Millero, 1985; Stumm and Morgan, 1996 etc.).
Stumm and Morgan (1996) compared Eq. (27) with fitted curves obtained from
empirical data attained from the oxidation of ferrous solutions with concentrations
lower than 0.5 mM/L (~28 mg Fe/l), at the pH range between ~ 4 and 6. In Fig. 5 it
can be seen that below pH ~4 the rate of oxidation is very low, and is fundamentally
independent of pH. Millero (1985) presented data for the rate of oxidation of Fe(II) in
solution which, among other observations, includes measurements at pH values higher
than 6 (Fig. 6). The results of Millero (1985) indicate that at pH values greater than ~8
the rate of ferrous oxidation is also independent of pH (these data was obtained with
seawater but the author extrapolated it also to freshwater). At pH values lower than
~8, the results of Millero (1985) are in agreement with Fig. 5. Thus the overall Fe(II)
oxidation rate versus pH plot shows a central region where the rate is strongly pH
dependent, flanked by regions on either side where the rate does not change as a
function of pH. Note that in freshwater above pH 9.0 the concentration of soluble
Fe(II) species is so low that it is very difficult to obtain reasonably accurate empirical
kinetic data (Morgan and Lahav, 2007).
34
Figure 5. Oxidation rate of Fe(II)
Figure 6. Rate constants for oxidation of
species as a function of pH ( PO2 = 0.20
soluble Fe(II) species in water as a function
atm). (Source: Morgan and Lahav)
of pH (Source: Millero,1985)
The trend apparent from these observations can be explained in terms of the
equilibrium chemistry of Fe(II) in aqueous solutions. The following rate equation
given by Millero (1985) separates the individual Fe(II) species present in aqueous
solution:
−
d[Fe(II)]
−
= (k o ([Fe 2+ ] + k 1[Fe(OH) + ] + k 2 [Fe(OH)02 (aq) ] + k 3 [Fe(OH)3 ])DO (28)
dt
Where: ko, k1, k2 and k3 are oxidation rate constants (time-1).
The value of dissolved oxygen concentration (DO) is used instead of PO2 since it is in
fact DO which participates in the oxidation reaction and the DO concentration will
vary for a given PO2 for a variety of reasons such as temperature, ionic strength and
the oxygen consumption rate of the solution (Millero et al., 1987).
Eq. (28) implies that in the course of soluble ferrous species oxidation, more than one
oxidation reaction occurs simultaneously (Stumm and Morgan 1996), e.g.
a
Fe 2+ → Fe 3+ + e −
b
Fe(OH ) → Fe(OH ) + e −
c
Fe(OH )2 → Fe(OH )2 + e −
+
0
2+
+
35
It implies also that the rate of oxidation depends on the concentration of each species
and its individual oxidation rate.
Fig. 7 is a log species – pH diagram for Fe(II) species in single phase aqueous
equilibrium (Total Fe(II) concentration = 10-3 M) and Fig. 8 is a log species – pH
diagram for Fe(II) for a two phase, aqueous – solid equilibrium, in the absence of
complexing agents for ferrous iron other than water (i.e. in the absence of Cl-, CO32-,
SO42-, etc.). Inspection of Fig. 8 reveals that in order to reach saturation at pH values
below ~8, the total Fe(II) concentration must be much higher than that used to
construct Fig. 5. In other words, below pH ~8 all Fe(II) solutions are effectively nonsaturated solutions and the equilibrium chemistry of Fe(II) solutions below ~pH 8 is
therefore better described by Fig. 7 (single phase aqueous equilibrium). For solutions
above ~pH 8 Fig. 8 (two phase aqueous-solid equilibrium) can be used to predict the
fraction of the ferrous concentration that is in the solid phase (Fe(OH)2(s) in this case)
(Morgan and Lahav, 2007). Once formed, the rate of oxidation of Fe(OH)2(s) has been
shown to be independent of pH, and to depend only on the rate of the introduction of
the oxidant (typically O2) into the water (Prasad and Ramasastry,1974 quoted by
Morgan and Lahav, 2007).
pH
pH
0
2
4
6
8
10
12
0
14
Fe(OH)3-
6
8
10
12
14
Fe2+
log species
-2
FeOH+
0
-9
Fe
(O
H
)2
log species
Fe2+
-6
-12
4
0
0
-3
2
Fe(OH)2(s)
-4
-8
-15
FeOH+
-6
-10
FeT
Fe(OH)20
Fe(OH)3-
Figure 7. Log species – pH diagram of
Figure 8. Aqueous – solid phase
soluble ferrous hydroxide species at
equilibrium for soluble ferrous hydroxide
infinite dilution (Morgan and Lahav,
species at infinite dilution (in the absence
2007)
of both inorganic carbon species and ion
pairing species such as Cl-, SO42-, PO43etc.) (Morgan and Lahav, 2007)
King et al. (1995) developed a detailed model of Fe(II) oxidation in homogenous
solution based on the Haber-Weiss mechanism. It is the most widely accepted
36
mechanism to describe the oxidation of Fe(II) by O2(aq) and has been used as the basis
for several subsequent studies of iron speciation and redox cycling in natural systems
(Santana-Casiano et al., 2006; Pham and Waite, 2008). The mechanism describes four
one-electron steps for the reduction of the terminal electron acceptor O2(aq):
A
Fe( II ) + O2 ( aq ) → Fe( III ) + O2• (−aq )
•
B
Fe( II ) + O2• (−aq ) + 2 H + → Fe( III ) + H 2 O2 (aq )
C
Fe( II ) + H 2 O2 ( aq ) → Fe( III ) + OH (•aq ) + OH −
D
Fe( II ) + OH (•aq ) → Fe( III ) + OH −
= radical
Combining sets of equations (a, b, c) with (A, B, C, D), Stumm and Morgan (1996)
showed that the free energy changes which occur during the oxidation of Fe(II) by
oxygen were more negative for the oxidation of Fe(OH)20(aq) than for either Fe(OH)+
and Fe2+. In all three reactions a, b, c step A was found to be endergonic but was least
endergonic for reaction c; and for all three reactions a, b, c steps B, C, D were all
found to be exergonic but in each case reaction c was the most exergonic. Thus by
explicitly making a kinetic argument from thermodynamic data, Stumm and Morgan
(1996) deduced that step A was the slowest step in each case (because it was the most
endergonic) and was therefore rate-limiting; and that step A was the fastest for
reaction c (because, of the three, it was the least endergonic). Similarly, they
concluded that since reaction c is most exergonic for steps B, C, D, all in all
Fe(OH)20(aq) produces the fastest oxidation sequence. This conclusion is backed up by
two other sources. Firstly, it has been shown that ‘hydrolyzed’ ferrous iron species are
more readily oxidized than non-hydrolysed ferrous species in the following order
Fe(OH)20(aq) >> Fe(OH)+ >> Fe2+, the reason presumably being that OH- ligands
donate electron density through both the σ and π systems to the reduced metal ion
which increases reducing power and stabilises the Fe3+ formed during the oxidation
(Luther, 1990 quoted by Stumm and Morgan, 1996). This hypothesis was
experimentally substantiated (in retrospect) by Millero (1985), who showed that the
rate constant for the oxidation of Fe(OH)20 is 5 orders of magnitude higher than the
rate constant for FeOH+, which, in turn, is 5 orders of magnitude greater than the rate
constant for Fe2+ reported by Lowson (1982). Secondly, Wehrli (1990) (quoted by
Morgan and Lahav, 2007) showed a linear free energy relationship between the free
37
energy of the reaction (log K) and the rate (log k) of the reaction for the oxidation of
Fe2+, FeOH+ and Fe(OH)20(aq).
Based on the studies of Lowson (1982) and Millero (1985) the following explicit rate
equation for soluble ferrous hydroxide species oxidation can be written (Eq. (29)).


− d[Fe(II)] 
6 ⋅ 10- 5 ⋅ Fe(II)T
1.7 ⋅ Fe(II)T
4.3 ⋅ 105 ⋅ Fe(II)T
=
+
+
2
3
+
2
K1K w K1K 2 K w K1K 2 K w
(H ) K 2 K w K 2 K 3K w
KK
(H + )
(H + )2
dt
1+
+
+
1+ 3 + w +
+
+
+ 2
 1 + (H + ) + (H + )2 + (H + )3
K1K w
(H )
(H )
(H ) K 2 K w K1K 2 K 2w




 × DO


(29)
Where: -dFe(II)/dt is the rate of ferrous oxidation in mole-Fe(II)/l/min and DO
corresponds to O2(aq) in equilibrium with PO2 =0.20 bar.
Each species is expressed explicitly as a function of the total ferrous iron
concentration (Fe(II)T), pH and stability constants. The effect of ionic strength and
temperature is determined correspondingly by adjusting the equilibrium constants by
the Davis equation (based on the Debye - Huckel approach) and the Van't Hoff
equation (Morgan and Lahav, 2007). Note that no rate constant has been reported for
the oxidation of Fe(OH)3- probably due to empirical limitations, however the
concentration of this species is exceptionally low at pH values below pH ~10 and
thus, from a practical standpoint, this rate constant is of a lesser importance. Also note
that the rate constants in Eq. (29) should probably be considered only with respect to
their order of magnitude, rather than as accurate quantities.
Having established the relative rates of oxidation of Fe(II) species one can now
rationally explain the oxidation rate curve seen in Fig. 5 and Fig. 6. It has already
been established that Fig. 7 applies only to Fe(II) solutions at pH below ~8.
Examination of Fig. 7 shows that in this pH range the concentrations of FeOH+ and
Fe(OH)20 rise steeply and linearly with pH. Since these species (especially Fe(OH)20)
are far more readily oxidized than Fe2+ is, this explains the pH-dependence of the
oxidation rate between pH ~4–8. Indeed, above pH ~5.5 both the first two terms on
the right hand side of Eq. (28) fall away (Stumm and Morgan 1996), which accounts
for the second order dependence of the rate law shown in Eq. (27) for this pH region.
Below pH ~4 the concentrations of FeOH+ and Fe(OH)20 are so low as to be
negligible and Fe2+ dominates. Since the oxidation of Fe2+ is independent of hydroxyl
groups, the oxidation rate is no longer pH dependent.
38
For pH values above ~8, all Fe(II) solutions of the order of at least 10-3 molar
concentration (under the assumption that ferrous complexes only with H2O) will be
saturated solutions. Fig. 8 shows the speciation of Fe(II) in two phase aqueous – solid
equilibrium. Here at pH ~8 although Fe2+ exceeds both FeOH+ and FeOH20 in
concentration, the latter two species are far more easily oxidized due to the presence
of the hydroxyl groups. As one moves from pH ~8 to the right in Fig. 8, both the Fe2+
and FeOH+ concentrations drop steeply whilst the FeOH20 concentration is constant.
As the pH increases above 8 the oxidation rate (Eq. (28)) therefore becomes
dominated by Fe(OH)20. Since this species is constant with respect to pH ‘throughout’
the range, the rate above pH ~8 becomes independent of pH. As explained earlier all
Fe(II) solutions of concentrations of the order of 10-3 M are unsaturated below pH ~8,
thus the Fe(OH)20 concentration is not constant below pH ~8 (as in Fig. 7 and not Fig.
8) (Morgan and Lahav, 2007).
6.3.2. Fe(II) oxidation by O2 in natural waters
In the presence of complexing or ion pairing agents other than H2O, as in natural
waters, more components should be added to Eq. (28). King (1998) suggested the
following oxidation rate:
−
d [Fe( II )]
= [O2 ] × [Fe( II )]× k app
dt
(30)
(
)
where k app = 4 k1α Fe 2 + + k 2α FeOH + + k 3α Fe (OH ) 0 + ...k nα n and αi is the fraction of each
2
Fe(II) species in solution, ki is the second-order rate constant for oxidation by oxygen,
and [Fe(II)] is the total or analytical Fe(II) concentration.
The factor 4 in Eq. (30) reflects the stoichiometry of Fe(II) oxidation by oxygen when
the first oxidation step (according to the Haber-Weiss mechanism) is rate limiting.
The oxidation of Fe(II) with O2 has been widely studied for many systems in marine
and atmospheric science and hydrometallurgy. Most experiments were conducted
with micro-molar and nano-molar concentrations of Fe(II), as found in aquatic
systems.
Carbonate is an important ligand in all natural waters and may form a variety of
complexes with both Fe(II) and Fe(III).
39
King (1998) conducted experiments in pure water with initial Fe(II) concentrations
between 1 to 5 micromolar. According to King's (1998) results the ferrous carbonate
complexes (FeCO30, Fe(CO3)22-, and Fe(CO3)(OH)-) dominate the speciation of Fe(II)
in natural waters containing greater than 1 mM carbonate alkalinity, where FeCO30 is
the most dominant species (see Fig. 9). In contrast, the Fe(CO3)22- complex is the most
kinetically active species for pH values above 6, while at pH values below 6.0, the
oxidation rate of Fe(II) is well described in terms of the Fe2+ and FeOH+ species (Fig.
10). Santana-Casiano et al. (2005) and Pham and Waite (2008) who investigated the
Fe(II) oxidation rate at nanomolar levels (0.025 - 0.250 µM) in natural waters,
attained the same tendency in their results.
Figure 9. Fe(II) speciation in pure water
Figure 10. Contribution of specific Fe(II)
with 2.3 mm of NaHCO3 and 0.03 atm of species in total Fe(II) oxidation rate by
PCO2 . (Source: King, 1998)
O2. Calculations are for pure water with
2.3 mM NaHCO3 at 25 °C. (Source: King,
1998)
In seawater, chloride and sulphate complexation with Fe2+ and increased ionic
strength are the significant processes responsible for reduced rates as compared to
pure water (Sung and Morgan, 1980; Lowson, 1982; Millero, 1985; King, 1998). The
complexation with Fe2+ decreases the Fe2+ fraction by a factor of 2 and
correspondingly shifts all other species by a comparable amount (King, 1998;
Santana-Casiano et al., 2006) (see Fig. 11 in comparison with Fig. 9). Since FeCl+ and
FeSO4 species have negligible contribution to the oxidation rate it explains the
40
observed rate in seawater, which is usually a factor of 100 times slower compared to
solutions of freshwater-like composition (Sung and Morgan, 1980).
Figure 11. Fe(II) speciation in 0.7 m of NaCl and 0.03 m of Na2SO4 with 2.3 mm of
NaHCO3 and 0.03 atm of PCO2 . (Source: King, 1998)
6.3.3. Fe(II) oxidation by O2 at low pH levels
The oxidation rate of Fe(II) by oxygen at low pH levels (pH < 2) has been examined
under various conditions. Table 6 presents results of studies conducted at near-room
temperatures (partially adopted from Lowson, 1982). In most studies the rate is not pH
dependent, as was explained in Section 1.3.1. If one compare these rates to the oxidation
rates in Table 5 one can see that the Fe(II) oxidation at pH<2 is very slow.
However, some possible catalysts are reported in literature. The following solids and
solutes were found to have a catalytic effect: in near neutral and neutral waters - alumina,
silica, bentonite and ferric hydroxide precipitates; in acid solutions - the surface catalysts
palladium, platinum, gold, and coconut charcoal and the solution catalyst Cu2+, Co2+,
PO42-, P2O72- and organic chelating agents (Stumm and Lee, 1961; Lowson, 1982;
Houben, 2004).
The anions PO42- or P2O72- increase the oxidation rate of Fe(II) by O2 due to complexing
of the ferric ion product (Cher and Davidsson, 1955; King and Davidsson, 1958). The
anions F- and Cl- show similar effects (King and Davidson, 1958; Rönnholm et al., 1999).
In general, the rate increases as the complexing affinity of the anion for ferric ion
increases (Huffman and Davidson, 1956; King and Davidson, 1958; Stumm and Morgan,
1996). Thus, at a given pH, the rate is found to decrease in the series pyrophosphate,
41
phosphate, chloride, sulfate and perchlorate. It is to be expected that the complexing
affinity for ferrous ion would be of the same order. The reported rate law is
− d (Fe( II ) )
dt
= 4k (Fe( II ) )PO2 for the first two media listed above and
dt
= 4k (Fe( II ) ) PO2 for the last three. It has been reported that the reaction is
− d (Fe( II ) )
2
also bimolecular with respect to Fe(II) and O2 in the presence of fluoride (Huffman and
Davidson, 1956).
Another possible catalyst is Cu2+. It is well known that addition of trace quantities of
Cu2+ can catalyze the oxidation of dissolved Fe(II) by O2 in acid to neutral pH (Cher and
Davidson, 1955; Huffman and Davidson, 1956; Stumm and Lee,1961; Lowson, 1982;
Chmielewski and Charewicz, 1984; Zhang et al., 2000).
In conclusion, the oxidation rate of ferrous iron by oxygen in natural waters is rapid due
to the presence of hydroxyl ions and the carbonate system, which form complexes with
Fe(II) that oxidize rapidly. At low pH levels the main species is the free ferrous ion that
has a very low oxidation rate. However, the anions in solution may have a significant
effect on the oxidation rate. In general, the rate increases as the complexing affinity of the
anion for ferric ion increases. Pyrophosphate and phosphate anions have the highest
acceleration effect. A catalytic effect by Cu2+ is also reported to be significant.
42
Table 6. Survey of reported Fe(II) oxidation rates under various conditions (Source: Lowson, 1982)
Maximal Rate
[M/s]*
Rate equation
2
kt х [Fe(II)] х pO2 +
8.68E-11
2+
2+
[Fe2+]
Rate constant
-5
-1 -1
-1
kt=1.5×10 M s atm
-3
-1 -1
39
kCu х [Fe ][Cu ]
kCu=7.6×10 M s
3.65E-10
k х [Fe2+]2 х pO2
2.78×10-6 M-1s-1atm-1
5.37E-6
1.23E-8
1.83E-8
3.36E-8
3.78E-8
4.68E-8
k1+k2 х [Cl-]
k х [Fe2+]2 х pO2
k х [Fe2+]2 х pO2
k = [H+]-0.23
k х [Fe2+]2 х pO2
k х [Fe2+]2 х pO2
k х [Fe2+]2 х pO2
1.90E-7
0.001 M
[media]
Media
0.5 M
1.10E-05
M
SO4
pH
T
[0c]
Reference
~1
30
Huffman and Davidson,
1956
2-
2+
Cu
1M
H2SO4
~0
30
k1=2.7×10-6, k2=8.9×10-7
6.5×10-7 M-1s-1atm-1
0.001-0.025
M
0.2 M
0.01-0.3 M
0-3 M
1N
HCl
HCl
<0
~0
30
30
Huffman and Davidson,
1956
Matseevskii, 1980
McBain, 1901
9.7×10-7 M-1s-1atm-1
0.01-0.3 M
1N
HClO4
~0
30
George, 1954
4.0×10-6 M-1s-1atm-1
2×10-6 M-1s-1atm-1
1.1×10-6 M-1s-1atm-1
0.15-0.2 M
0.01-0.3 M
0.02-0.45 M
1M
1N
1M
H2SO4
HClO4
H2SO4
~0
~0
~0
30
30
30
kх[Fe2+]2х pO2 х[H+]-0.25
5.1×10-7 M-.0.75s-1atm-1
0.01-1.0 M
1N
H2SO4
~1
30
kх[Fe2+]х pO2 х[H3PO4]2
k1х[Fe2+]хpO2 х[H3PO4]2
+
7.88E-6
2+
k2х[Fe ]хpO2х[H4P2O7]
kх[Fe2+]хpO2хf{HCl}
kх[Fe2+]хpO2хf{HCl}
*
mg Fe(II)/l/h = 2.0х108 M/s
1.25×10-3 M-2s-1atm-1
0.005-0.02 M
1N
H3PO4
H3PO4
+
H4P2O7
HCl
HCl
~2
30
Lamb and Elder, 1931
Chernaya, 1980
McBain, 1901
Mathews and Robins,
1972
Cher and Davidson, 1955
1-2
30
King and Davidson, 1958
<1
<1
18
45
Posner, 1953
Iwai et al, 1979
5.83E-7
k1=1.08×10-3 M-2s-1atm-1
k2=2.13×10-3 M-1s-1atm-1
1.3×10-3 s-1atm-1
3.1×10-4 s-1atm-1
0.005-0.02 M
0.055 N
0.2 M
43
0.2-0.4 M
0.2-0.8 M
8N
6N
6.4. Electrochemical oxidation of Fe(II)
The use of electrolysis to oxidize ferrous iron has been investigated mainly in
association with treatment processes for acid mine drainage (AMD) (Bunce et al.,
2001) and disinfection of drinking water (Khelifa et al., 2004).
The regeneration of ferric iron in an electrolytic cell can be achieved by direct or
indirect oxidation. Direct oxidation is referred as the oxidation of Fe(II) to Fe(III) on
the anode surface, whereas indirect oxidation is accomplished by producing a strong
oxidizing agent through electrolysis that is subsequently used to oxidize the Fe(II).
6.4.1. Direct oxidation of Fe(II)
The direct electro-oxidation process is based on oxidation of Fe(II) to Fe(III) on the
anode surface (Eq. (31)) and evolution of hydrogen gas on the cathode surface (Eq.
(32)) (Bunce et al., 2001).
Fe 3+ + e − → Fe 2+
anode
EH0=0.769 V
(31)
2 H + + 2e − → H 2
cathode
EH0=0.092 V
(32)
A competing (unwanted) reaction that invariably occurs on the anode is the oxidation
of water to dissolved oxygen:
1
2
O 2 + 2 H + + 2e − → H 2 O
anode EH0=1.226 V
(33)
Additionally, Ferric ions formed at the anode and those that are initially present in the
solution can be back reduced at the cathode (Bisang, 2000; Bunce et al., 2001).
6.4.2. Indirect oxidation of Fe(II)
The electro-generation of strong oxidizing agents, especially chlorine and/or
hypochlorite for the removal of low concentrations of organic matter is a process that
is known for many years (Alvarez-Gallegos and Pletcher, 1998). When elemental
chlorine (Cl2(g)) is introduced into the water at natural pH it rapidly reacts with H2O
(Eq. (34)) to form hypochlorous acid:
Cl 2 ( g ) + H 2 O → HOCl + H + + Cl −
pK=3.40
(34)
Hypochlorous acid exists in solution in equilibrium with hypochlorite ions (OCl-)
according to pH (pKa= 7.46), as shown in Eq. (35).
HOCl ↔ OCl − + H + pK a = 7.46
44
pKa = 7.46
(35)
These three oxidizing forms are denoted Active chlorine or AC in short.
At the pH range 3-10.5 the hydrolysis rate (Eq. (34)) is rapid and almost constant. At
pH values above 10.5 the rate sharply increases with an increase in pH (Spalding,
1962). However, at strongly acidic solutions (pH<3) chlorine hydrolysis is negligible
(Spalding, 1962) and the dominant species in the aqueous phase is Cl2(aq).
Since transporting, storing and handling of liquid chlorine are dangerous and
ecologically unsafe (Kelsall, 1984; Krstajić et al., 1987) investigation and
development of local, on-site processes for active chlorine production has gained
increasing interest. The major process for an on-site AC production is electrolytic
conversion of chloride ions in sodium chloride (NaCl) solutions into desired
hypochlorite or chlorine gas as in Eq. (36).
Cl 2 ( g ) + 2e − → 2Cl − anode EH0=1.392 V
(36)
The oxidation of ferrous iron by AC has been investigated mainly in the context of
disinfection for drinking water (Khelifa et al., 2004). The products of Fe(II) oxidation
are Fe(III) and chloride ions:
or
Cl 2 + 2 Fe 2+ → 2Cl − + 2 Fe 3+
(37)
HOCl + 2 Fe 2+ + H + → Cl − + 2 Fe 3+ + H 2 O
(38)
6.4.2.1. Competing reactions
Several competing reactions for the electro-generation of active chlorine have been
reported, contributing to cell efficiency reduction and byproduct formation, (Krstajić
et al., 1991; Rudolf et al., 1995):
At the anode hypochlorite may be oxidized to chlorate according to Eq. (39).
−
6OCl − + 3H 2 O → 2ClO3 + 4Cl − + 6 H + + 12 O2 + 6e −
anode
EH0=1.446 V (39)
Another competing reaction at the anode is oxygen formation (Eq. (33)).
Hypochlorite may be also reduced at the cathode to the chloride ion form as described
in Eq. (40) or chemically transformed into chlorate or chloride according to Eq. (41)
and Eq. (42) respectively:
OCl − + H 2 O + 2e − → Cl − + 2OH −
cathode
EH0=1.481 V (40)
−
2 HOCl + OCl − → ClO3 + 2Cl − + 2 H +
(41)
2OCl − → 2Cl − + O2
(42)
45
The ratio between the anode surface area to the cathode surface area (Sa:Sc) may
affect the rate of chlorate production on the anode and its reduction on the cathode.
Khelifa et al., (2004) reported that the optimum Sa/Sc value in their experimental
system was 1.33 for anode made of Ti/TiO2 and cathode made of Ti.
Nevertheless, the rate of chlorate formation according to Eq. (41) depends on the
temperature and is higher at elevated temperatures. At relatively low temperatures
(about 20-25 °C) chlorate formation is slow (Grinberg et al., 2001). The pH of
solution also affects chlorate formation. First of all Cl2 hydrolysis to HOCl is
relatively slow below pH3 (see above), and since hypochlorite is the reactant for
chlorate production - chlorate formation should be minor below pH3. Secondly, the
main reaction on the cathode around pH2 was reported to be H2 evolution (AlvarezGallegos and Pletcher, 1998; Bisang, 2000) (Eq. (32)).
6.4.2.2. Oxidation of Fe(II) by Cl2
The following mechanism was proposed by Gilliland et al. (1958) and confirmed by
Hikita et al. (1975):
step I: Cl 2 + Fe 2+ ↔ Cl 2 ⋅ Fe 2+
(43)
step II: Cl 2 ⋅ Fe 2+ + Fe 2+ ↔ 2Cl − + 2 Fe 3+
(44)
The first step is an irreversible second-order reaction (first order with respect to both
ferrous iron concentration and chlorine concentration) and is the rate controlling step
of the overall reaction. The second step is expected to be second order (Gilliland et
al., 1958).
The information on oxidation of Fe(II) by Cl2 is scarce. Yet, the electro-oxidation of
chloride ions to active chlorine is prevalent in literature.
6.4.3. Electrolytic production processes of active chlorine
The production of active chlorine depends on various operational parameters; some of
these parameters are discussed hereunder.
6.4.3.1. Divided or undivided cell
Electrolytic cells appear in the form of divided or undivided cells. In an undivided
cell, the cathode and anode electrodes are merged in the same solution, without any
46
separation of the electrolyte solution. A divided cell, on the contrary, is divided into
compartments – one for each electrode, which make two different electrolyte
solutions: an anolyte solution for the anode and a catholyte solution for the cathode
(Bunce et al., 2001).
Divided cells produce more concentrated hypochlorite solutions and have higher
current efficiency for hypochlorite production as compared with undivided cell
processes (current efficiency is the ratio of the electrochemical equivalent current
density for a specific reaction to the total applied current density). The downside of
this method is that electrolysis in membrane-divided-cells for hypochlorite and
chlorine production is susceptible to failure due to clogging and formation of
precipitates on the membrane surface (Bisang, 2000). This affects membrane
permeability and, consequently, overall process efficiency.
An undivided cell is prone to current efficiency loss due to the unwanted reactions
and byproducts formation, out of which chlorate formation is probably the most
problematic. However, as mentioned earlier, working at room temperature and at low
pH decreases remarkably the production of chlorate.
6.4.3.2. Batch or flow-through mode
Electrolytic cells can be operated in batch mode where the solution is continuously
mixed during electrolysis (Krstajić et al., 1987; Khelifa et al., 2004).
Another approach is flow-through electrolysis where two vessels are used, one as the
solution carrier and another as the electrolytic cell. The electrolyte is recycled from
the holding vessel to the electrolytic cell (Rudolf et al., 1995; Kraft et al., 1999;
Grinberg et al., 2001). The flow-through cell can be operated as a single pass mode
(Kelsall, 1984; Grinberg et al., 2001) or in a recirculation mode (Robertson et al.,
1983; Rudlof et al., 1995). Flow-through systems are favorable because the forced
convection prevents local overheating of the electrolyte near the electrode surface
(Grinberg et al., 2001), reduces gas locking of the cathode, provides similar current
density distribution across the electrodes area and prevents the precipitation on the
electrodes (Kelsall, 1984).
Other operational parameters to be considered are: flow-through velocity, chloride
concentration, anode and cathode materials, interelectrode gap (i.e., the distance
between the anode and the cathode) and current density.
47
In Table 7 the conditions and results of three experiments for hypochlorite production,
all in undivided cells, are summarized. The conditions of the experiments presented in
Table 7 vary, but one can see that chlorate formation rates of 8 g/h can be achieved,
which means that similar rates are possible for Fe(II) oxidation.
Table 7. Conditions and results of experiments on chloride electrochemical oxidation
in an undivided cell
Rudolf et al., 1995
Cell type
Flow-through
rate
Initial NaCl
concentration
Electrolyte
temperature
and pH
Grinberg et al., 2001
Khelifa et al., 2004
Flow-through,
Flow-through, single
recirculating
pass
5.16 l/min
0.20 l/min
-
15 and 30 g/l
8.9 g/l
3M
25°C
initial pH=6
20°C
batch
20°C
pH>8
Titanium covered
Anode
material
with 30 mol%
Ti/Pt
Ti/RuO2
Stainless steel
Ti/Pt
Titanium
1
0.92
1.33
4 mm
0.5 mm
5 mm
0.12 kA/m2
3.5 kA/m2
RuO2 and 70 mol%
TiO2
Cathode
material
Sa/Sc ratio
Interelectrode
gap
Current
1.059 – 4.237
densities
kA/m2
sodium
Calculated
hypochlorite: 7.8-
oxidation rate
10.0 g/h / chlorate:
0.3-0.9 g/h
active chlorine: 8.2
sodium hypochlorite:
g/h / the only active
11.0 g/h / chlorate:
chlorine species
<0.2 g/h
detected was
hypochlorite
48
In summary, it is possible to use an electrolytic process to oxidize Fe(II), directly or
indirectly (i.e., in the presence of chlorides). The avoidance of competing reactions
should be considered. Many operational parameters can be optimized in order to
achieve a high oxidation rate of Fe(II).
In the current work the oxidation of Fe(II) by O2 in the presence of catalysts and
electrochemical oxidation of Fe(II) were investigated as part the development of an
LRSR process at low pH.
49
7. Hypothesis and objectives
7.1. Research hypothesis
It is possible to develop an LRSR process that is operated at low pH (pH<1.5), under
such conditions where: (1) precipitation is insignificant and does not impede the
process, (2) reactive-absorption of H2S(g) is efficient and (3) Fe(III) regeneration (i.e.
Fe(II) oxidation) rate is equal or higher than the rate of H2S(aq) oxidation by Fe(III).
Two process options were investigated with respect to the research hypothesis:
LRSR process at pH1.0 with catalytic oxidation of Fe(II) by atmospheric O2
This option is based on the use of Cu2+ and H3PO4 which have been reported in the
literature to accelerate the Fe(II) oxidation rate at low pH. A bubble column reactor is
filled with solution that contains high concentration of Fe(III) (9 g Fe/l) and copper
and/or phosphate. A gaseous stream contaminated with H2S(g) is introduced into the
reactor. The concentration of Fe(III) is maintained high in order to obtain a high
reactive absorption efficiency. The Fe(II), generated as a result of H2S oxidation by
Fe(III), is oxidized back to Fe(III) by O2, which is introduced into the reactor with the
gaseous stream. If the O2 content in the gaseous stream is low, air or oxygen should
be supplied to the reactor. The Fe(II) oxidation rate in the presence of copper and/or
phosphate should be high enough in order to maintain the Fe(III) concentration high
and steady.
LRSR process at pH1.0 with electrochemical oxidation of Fe(II)
This option is based on the use of an electrolytic cell to accelerate the Fe(II) oxidation
rate. A bubble column reactor is filled with solution that contains a high concentration
of Fe(III) (9 g Fe/l). The oxidation of the Fe(II) can be preformed directly or
indirectly. In the case of indirect oxidation chloride ions are also added to the
solution. A gaseous stream contaminated with H2S(g) is introduced into the reactor.
When the concentration of Fe(II) generated exceeds a certain level (2 g Fe/l), a
portion of the solution is recycled into a flow-through electrolytic cell. In the direct
oxidation option the Fe(II) is oxidized on the anode surface, while in indirect
oxidation option Cl- is oxidized on the anode surface to Cl2 that oxidizes the Fe(II).
When the concentration of Fe(II) goes down below a certain level (1 g Fe/l) the
circulation into the electrolytic cell is stopped. Such an operation enables maintaining
a relatively constant and high Fe(III) concentration, while the fact that Fe(II)
50
oxidation is not operated to very low Fe(II) concentrations has a kinetically
advantage.
The use of a bubble column reactor has the advantage of a high contact area between
the gas and the solution which suppose to reduce the size of the scrubbing reactor.
Apart from that, the type of the reactor that is used is of no importance to the process.
7.2. Research objectives
1) Assessment of the potential of long-term formation of Fe(III) precipitates from the
H2S reactive-absorption solution as a function of pH and the composition of
solution.
2) Investigation of the effect of pH and H3PO4, Cu(II) and Cl- concentrations on the
reactive-absorption efficiency of H2S(g).
3) Investigation of the effect of Fe(III), H3PO4 and Cu(II) concentrations on Fe(II)
oxidation by O2 at pH1.0.
4) Investigation of the effect of Cl- concentration and current density on Fe(II)
electrochemical oxidation.
5) Determining the most advantageous conditions (or range of conditions) for
catalytic oxidation of Fe(II) at pH1.0, with respect to H2S(g) removal by the LRSR
process, and evaluation of the Fe(II) oxidation rate.
6) Determining the most advantageous conditions (or range of conditions) for
electrochemical oxidation of Fe(II) at pH1.0, with respect to H2S(g) removal by the
LRSR process, and evaluation of the Fe(II) oxidation rate.
51
8. Materials and Methods
8.1. Chemicals
All reagents were analytical grade if not stated otherwise.
Table 8. Description of the reagents used in experiments
Name of chemical
Formula
Acetic acid glacial
CH3COOH
Ammonium acetate
CH3COONH4
Ammonium iron(II)
sulfate 6-hydrate (FAS)
Copper sulfate
Manufacture
Remarks
BIO-LAB ltd.
99.8%
LOBA CHEMIE
(NH4)2Fe(SO4)2•6H2O Merck
CuSO4
Merck
Hydrochloric acid
HCl
Frutarom
Hydrogen sulphide
H2 S
BOC
Iron(II) sulphate 7hydrate
Iron(III) sulfate
exsiccated
FeSO4•7H2O
Fe2(SO4)3
Sigma-Aldrich
C.P.
85%
H3PO4
Carlo Erba
Sodium chloride
NaCl
Frutarom
Sodium fluoride
NaF
Riedel de Haen
Sodium hydroxide
NaOH
Frutarom
Sulfuric acid
H2SO4
Frutarom
monohydrate
5-Sulphosalicylic acid 2hydrate
Nitric acid
Mercury thiocyanate
C12H8N2•H2O
C7H6O6S•2H2O
Min. 98%
Fluka
Riedel de Haen
HNO3
Merck
C2HgN2S2
Fluka
52
Background
Carlo Erba
ortho-Phosphoric acid
1,10-Phenanthroline
500 ppm, N2
65%
8.2. Analytical equipment
1) Filter paper, 0.45 µm, White 47 mm Gridded, Sterile, Millipore
2) (ICP) emission spectrometer, Optima 3000 DV, Perkin Elmer
3) Kitagawa gas detector tube system for H2S(g) (1-60 ppm, 0.5-40 ppm and 0.2-6.0
pH meter, Metrohm 827 pH lab, with Metrohm 6.0228.010 glass electrode
4) Linear shaker with the water bath, BT-350, MRC Laboratory equipment
5) MINEQL+ software version 4.5 (Envrionmental Research Software, Hallowell,
ME)
6) QuickChem 8500 (LACHAT instruments, USA) for ammonia detection
7) Spectrophotometer, Genesys 10, Spectronics
8) Syringe Driven Filter Unit, 0.22 µm. Millex-GV, Millipore
9) ppm), Komyo Rikagaku Kogyo K.K.
10) Tubes gas dispersion pyrex fritted disc 40-60 microns
11) 0-20A DC power supply device (HY3020, Aviv Energy tech. Ltd, Israel)
8.3. Experimental
All experiments were conducted at room temperature (25 ±1 °C); pH adjustments
were done by the addition of concentrated (98%) sulfuric acid or NaOH 5N.
The applicability of each one of the proposed processes was investigated according to
the following parameters: (1) H2S(g) reactive-absorption efficiency, (2) practical
precipitation potential of the working solution and (3) Fe(II) oxidation rate.
8.3.1. H2S(g) reactive-absorption experiments
The efficiency of H2S(g) reactive-absorption into acidic solutions containing Fe(III)
was studied using a bubble column reactor. Mixtures of hydrogen sulfide and air were
bubbled through a bubble column reactor filled with acidic ferric sulfate solution. The
flow rates of both the air and the hydrogen sulfide gas were measured and controlled.
The total flow rate was kept low enough in order to maintain a homogeneous bubble
flow. The desired concentrations of H2S and mixed air flow rates were achieved by
controlling the individual flow rates of the N2-H2S mixture and air. The concentration
of H2S in the experiments did not exceed 80 ppm, mainly for reasons of safety. The
53
H2S(g) concentrations at the inflow and the outflow of the reactor were measured by
Kitagawa gas detector tubes designed for H2S(g). Each measurement was repeated two
or three times.
8.3.2. Determination of practical precipitation potential of the working
solutions
Solutions with different compositions were left for several weeks to months in order
to follow the formation of precipitates. Samples from solutions were taken and filtered
every few weeks and the concentration of several dissolved components (e.g., iron,
copper) was measured by ICP analysis or via a spectrophotometric method. Samples
of the solid phase were also taken few times. The samples were dried in an oven and
then dissolved in order to determine the presence of part of the components and their
concentrations ratio.
8.3.3. Determination of Fe(II) oxidation rate
8.3.3.1. Catalytic oxidation of Fe(II)
Batch experiments were conducted with initial ferrous iron concentration of 1000 mg
Fe/l (from (NH4)2Fe(SO4)2•6H2O) in 500 ml bottles. Air was bubbled to supply
oxygen through gas dispersion tubes. Fig. 12 is a photo of the system that was used
for the catalytic oxidation experiments. The dissolved oxygen concentration was
measured and verified to be the saturation concentration. Every few hours the bottles
were weighed and refilled to the initial weight with diluted H2SO4 solution (~0.1 N).
Samples from solutions were taken and diluted in diluted H2SO4 solution in order to
determine the dissolved Fe(II) concentration. The concentration of Fe(II) was
measured by the modified Phenanthroline method developed by Herrera et al. (1989).
Each sample was measured three times, and if the standard deviation of the three
samples was less than 5% the average value of the three was taken. Otherwise only
two replicates were considered.
54
Figure 12. A photo of the experimental system for catalytic Fe(II) oxidation
8.3.3.2. Electrochemical oxidation of Fe(II)
The electrochemical laboratory system setup that was used is shown in Fig. 13. A
flow-through electrolytic cell with parallel rectangular electrodes was used. The
electrolyte was recycled between the flow-through cell and the electrolyte holding
vessel by a peristaltic pump. A PVC pipe reactor, (internal diameter = 5.71 cm) was
used as a holding vessel to provide the Cl2(g) formed on the anode a longer retention
time inside the system. A 0-20A DC power supply device was used in order to
provide the electrical current. Three cathode dimensions were used: 0.72×9.15 cm2,
3.0×9.14 cm2 and 5.1×9.15 cm2. Anode dimensions were 5.1×9.15 cm2 in all
experiments.
Electrode materials, inter-electrode gap and flow-through velocity were adopted from
literature related to the electrochemical hypochlorite production, as discussed
hereunder.
The most effective anode material for hypochlorite production was reported to be
titanium coated by ruthenium oxide (Ti/RuO2) due to its low over-potential for
chloride oxidation and mechanical and chemical stability (Ponzano, 2007). Cathode
material may affect the hypochlorite back reduction rate and hydrogen evolution rate
(Khlaifa et al., 2004). Ponzano (2007) recommended on titanium coated with iridium
55
oxide and plain titanium as good cathode materials, while graphite, stainless steel and
nickel were found less suitable. In the current study titanium electrodes were used.
Khelifa et al., (2004) studying a batch electrolyzer for hypochlorite production and
reported the optimal interelectrode gap to be <8 mm. In other studies the
interelectrode gap used was 4 mm or shorter (Kelsall, 1984; Robertson et al., 1983;
Rudolf et al., 1995). An inter-electrode gap of 3.46 mm was used in the current work.
The oxidation of hypochlorite to chlorate on the anode depends on the hypochlorite
concentration in the vicinity of the anode (Grinberg et al., 2001). Thus the flowthrough velocity may affect the rate of chlorate production. The flow-through
velocities that were found in literature were 26.275 cm/s (Rudolf et al., 1995) and
0.27-1.08 cm/s (Grinberg et al., 2001). The electrolyte recirculation rate between the
electrolysis cell and the holding vessel in the current work was 1.72 l/min and
consequently the flow velocity in the electrolysis cell was 16.24 cm/s.
Figure 13. A scheme (on the left) and a photo of the laboratory system for Fe(II)
electrochemical oxidation experiments. 1- Electrolyte holding (Cl2 absorption) vessel;
2- Flow-through electrolysis cell; 3- DC power supply; 4- Peristaltic pump; 5Chlorine trap.
In the direct oxidation experiments (Cl-=0) the concentration of dissolved Fe(II) was
detected by measuring the oxidation-reduction potential (ORP) at predetermined time
56
intervals. In the indirect oxidation experiments the concentration of dissolved Fe(II)
was determined using the modified Phenanthroline method (Herrera et al., 1989).
All the gas that was formed during the electrolysis was passed through a measuring
glass (250 ml) filled with 1M NaOH solution. At this alkaline solution Cl2(g) that
formed on the anode surface and did not react with Fe(II) was hydrolyzed according
to Eq. (34). Consequently, the determination of the chloride ions concentration in the
trap solution was used to assess the chlorine mass that potentially escaped the system
during the operation.
8.4. Analytical methods
8.4.1. Determination of the total dissolved iron, phosphate and copper
concentrations
The concentrations of the total dissolved iron, phosphate and copper were analyzed by
inductively coupled plasma (ICP) emission spectrometry, Optima 3000 DV, Perkin
Elmer.
8.4.2. Determination of dissolved ferrous iron concentration
The concentration of dissolved ferrous iron was measured spectrophotometrically
according to the modified phenanthroline method (Herrera et al., 1989). The method
is highly accurate (below 5% error) in the presence of ferric iron if the Fe(II)
percentage is higher than 5% of the total dissolved iron. Samples were diluted with
diluted H2SO4 solution (pH1.0) in order to lower the Fe(III) concentration below 200
mg Fe/l.
Two modifications were made to the method: the absorbance of a sample was read 10
minutes after the addition of the colorimetric reagent instead of 5 minutes and diluted
H2SO4 solution (pH1.0) was used instead of distilled water.
57
8.4.3. Determination of dissolved ferric iron concentration
The concentration of dissolved ferric iron was measured via a spectrophotometric
method using Sulphosalicylic acid (Zolotov, 2001). Samples were diluted with diluted
H2SO4 solution (pH1.0) in order to lower the Fe(III) concentration below 20 mg Fe/l.
8.4.4. Determination of chloride concentration
The concentration of chlorides was measured via a spectrophotometric method using
mercury thiocyanate and iron(III) alum (Yoshinaga and Ohta, 1990). Yoshinaga and
Ohta (1990) reported that higher accuracy was obtained with higher concentrations of
mercury thiocyanate and FAS reagents. Thus, the concentrations used in current work
were the highest reported by Yoshinaga and Ohta (1990), i.e., 9 g/l mercury
thiocyanate and 300 g/l of FAS.
The method is reliable also in highly basic solutions (1M NaOH), although a specific
calibration curve should be made.
8.4.5. Determination of ammonia concentration
The concentration of ammonia was determined by flow injection analysis of
QuickChem (method 10-107-06-2-A). The ammonia concentration is determined by a
spectrophotometric measurement, based on the color that is produced when ammonia
is heated with salicylate and hypochlorite in an alkaline phosphate buffer.
8.4.6. Analysis of species distribution by the MINEQL+ software
Input data - components of solution, total concentrations of the components and pH.
Ionic strength was calculated during the runs of the software.
Output data – concentrations of each one of the species of each component. Data on
precipitation potential of all possible solids is also given.
It should be mentioned that MINEQL+ analysis is based on thermodynamic data and
kinetics are not considered. In addition, ionic strength corrections in MINEQL+ are
calculated according to the Debye-Huckel equation, which means that the accuracy of
thermodynamic constants is lower at ionic strengths exceeding 0.5 M. In the current
58
study the ionic strength of most of the solutions exceeded 0.5M, thus results should be
considered to be semi-qualitative rather than quantitative.
59
9. Results and Discussion
The Results chapter is divided in two according to the two processes that were
investigated in the work: LRSR process at pH1.0 with catalytic oxidation by O2 of
Fe(II) and LRSR process at pH1.0 with electrochemical oxidation of Fe(II). In each
section the results of the following three subjects are presented: (1) reactiveabsorption efficiency of H2S(g), (2) practical precipitation potential of Fe(III) species
in the working solution and (3) Fe(II) oxidation rate, all under various operational
conditions.
9.1. Catalytic oxidation
The compositions of the solutions that were tested in the catalytic oxidation
experiments were:
Fe(III) concentrations: 0, 81, 161 and 322 mM = 0, 4.5, 9.0 and 18.0 g Fe(III)/l
Total-phosphate concentrations: 0 to 867 mM = 0 to 26.7 g P/l
Cu(II) concentration: 0 to 8 mM = 0 to 0.5 g Cu(II)/l.
The initial Fe(II) concentration was 18 mM (= 1 g Fe(II)/l) in all experiments.
4.2.4. H2S(g) reactive-absorption efficiency
The reactive-absorption efficiency was tested once in a solution with the following
composition: 18 mM Fe(II), 81 mM Fe(III), 220 mM H3PO4 and 2 mM Cu(II) at
(
)
3
3
and it contained around 80 ppm
pH1.22. The flux of gas was 1.25 mair
/ min / m solution
of H2S(g). The reactor had an internal diameter of 5.71 cm, and air was sparged by a
gas dispersion tube.
In less than one hour a colloid-type dark precipitate appeared in the solution. The
precipitate was identified as the mineral Covellite (CuS). Covellite is described as thin
crusts or as deep blue-black powdery or sooty masses (Dunn and Muzenda, 2001).
The structure of CuS covellite is best represented as Cu(I)4(S-I)4Cu(II)2S(-II)2
(Madarász et al., 2001). The precipitation of CuS was corroborated also by the
measurements of dissolved copper concentration. As shown in Fig. 14 the
concentration of dissolved copper decreased linearly with time, which confirmed the
CuS precipitation observation.
The reactive-absorption efficiency of H2S(g) with in the experiment time is also
presented in Fig. 14. The reactive-absorption efficiency was high at the beginning of
60
the experiment (80-85%), and then decreased. This is probably due to the decrease in
the availability of Cu(II), which precipitated with S2- to form CuS. This means that the
oxidation of H2S(aq) in the tested solution was relatively slow and that precipitation of
CuS was the main mechanism by which H2S(g) was removed.
Zhang and Tong (2006) claimed that once CuS forms, it is rapidly oxidized by Fe(III)
to S0 (see1.2.4.2). This is in contrast with the results presented here. It is possible that
the higher concentrations of Fe(III), Fe(II) and Cu(II) and the different concentrations
ratio, that were applied by Zhang and Tong (2006) (see Subsection 1.2.4.2), are the
reason for this difference in results. However, the very high concentrations that were
suggested by Zhang and Tong (2006) are problematic for blow-down reasons. The
solution contains high concentration of copper, whose release to the environment is
limited. In addition, Zhang and Tong (2006) reported that no precipitation (other than
elemental sulfur) was observed after 4 hours of experiment. Yet, it is very unlikely
that in such a concentrated solution precipitation will not occur in long-term
[%]
H2S(g) reactive-absorption efficiency
operations (which are much longer than the 4 hours they reported on).
0.90
1.6
0.85
1.4
0.80
1.2
0.75
1.0
0.70
0.8
0.65
0.6
y = -0.2352x + 1.4287
R2 = 0.993
0.60
0.4
0.55
0.2
0.50
0.0
0
1
2
3
Time [h]
4
5
6
Figure 14. Change in H2S(g) reactive-absorption efficiency (◊) and dissolved copper
concentration (х) with time.
61
In conclusion, the addition of Cu(II) into a solution that is used for an LRSR process
is problematic because at the conditions tested, because Cu2+ apparently precipitates
with S2- to form CuS.
The addition of H3PO4 as a catalyst was also found problematic, as discussed in the
following section.
4.2.5. Practical precipitation potential in the working solution
Precipitation was observed in all Fe(II) oxidation experiments (Section 4.1.3) in
which phosphate concentration in solution was 0.16 M (4.6 g P/l) or higher. It took
between a few hours to a few days for the precipitates to become visible. Two types of
precipitates were observed: the first was bright-yellowish and powder like; and the
second bright-purplish and crust like (Fig. 15). Samples from 5 different bottles were
filtered (after standing untouched for 170 days) and the solid phase was dried in an
oven at 60°C for 24 hours. The dried phase was then dissolved in highly acid solution
(HCl 1N) and the samples were taken for ICP analysis, in order to quantify the
content of iron, phosphate and copper in the solid. In all 5 samples the concentration
of copper was below the detection limit. The concentrations of total-phosphate (PT)
and total-iron (FeT) found in the solid phase are presented in Table 9. In all 5 samples
the molar ratio of phosphate to iron was found to be between 1 and 2.
The ICP analysis of the solutions from other Fe(II) oxidation experiments showed a
clear decrease in the concentration of phosphate and iron, relative to the initial
concentrations, in most of the samples. A decrease in ammonia concentration was also
observed, probably due to physical adsorption of ammonia on the solid surface.
Figure 15. A photo of precipitates from catalytic oxidation experiments after 170 days
62
Table 9. Composition of precipitates in solutions from the catalytic oxidation
experiments (after 170 days)
Total phosphate and total iron concentrations
Sample
FeT
FeT
Color
PT [mg/l]
2
yellow
9.8 ± 0.6
16.5 ± 0.1
0.32
0.30
1.1
3
yellow
55.7 ± 0.7
75.0 ± 0.2
1.80
1.34
1.3
4
purple
16.9 ± 0.1
26.2 ± 0.8
0.54
0.47
1.2
5
purple
16.8 ± 0.2
25.6 ± 0.6
0.54
0.46
1.2
6
purple
26.5 ± 1.6
25.8 ± 0.0
0.86
0.46
1.9
number
[mg/l]
PT [mM]
PT/FeT ratio
[mM]
M/M
The following solution was analyzed by MINEQL+ in order to evaluate its
thermodynamic precipitation potential and compare it to the observed results: 18 mM
Fe(II), 81 mM Fe(III), 59 mM H3PO4 and 0.8 mM Cu(II) at pH1.0. The precipitation
potential was found positive for the minerals Strengite (FePO4•2H2O) and H-Jarosite
(HFe3(OH)6(SO4)2). In fact, 99% of total-H3PO4 was expected to precipitate as
strengite at steady state, and the rest of the Fe(III) was projected to precipitate as
jarosite. This is equivalent to 20.6 g of solids for 1 l solution, which is much higher
than the amounts of solid that were visually observed to precipitate. In the actual
experiments, which lasted for 3 months it was assumed that at pH conditions as low as
pH1.0 the kinetics of precipitation would be very slow and thus solid precipitation
would appear after a very long time (in the order of months). If this was indeed the
case, the loss of Fe(III) due to precipitation would be tolerable, from an engineering
perspective.
The assumption of slow kinetics at this pH was supported by Cher and Davidson
(1955), who did not report on any precipitation in their experiments, which were
conducted at pH 1.1-1.7 and with PT between 0.402 and 1.0 M. However, in their
experiments the initial Fe(II) concentration was 10 mM and no Fe(III) was added,
which may explain the fact that precipitation was not observed: even if all the Fe(II)
was oxidized to Fe(III), the maximal precipitation possible is of 0.56 g Fe per 1 l of
solution..
63
All in all, the addition of the potential catalysts Cu2+ and H3PO4 to an LRSR solution
was found to be inapplicable in the range of conditions that were tested in the work.
Nevertheless, the investigation of the effect of Cu2+ and H3PO4 on Fe(II) oxidation by
O2 at low pH was continued because of its scientific value.
4.2.6. Fe(II) oxidation rate in the catalytic oxidation experiments
First, the effect of phosphate or copper on the modified phananthroline method (Herrera
et al., 1989) was assessed. Calibration curves were prepared with different concentrations
of H3PO4 and Cu(II). The different calibration curves that were attained were similar to
each other and no distinguishable effect of H3PO4 and Cu(II) was observed.
Consequently, a general calibration equation was chosen to represent all the solutions, as
follows:
y = 122.4 x − 1.902
(45)
where y is Fe(II) concentration in mg Fe/l and x is the absorbance in %.
All experiments were conducted for 4 to 20 hours. In each experiment the concentration
of Fe(II) was measured 6 to 12 times. Examples of the results of 4 different experiments
are presented in Fig. 16.
Fe(II) concentration [mg Fe/l]
1000
900
Fe(III)=81mM,
Cu(II)=0,
P=588mM
Fe(III)=161mM,
Cu(II)=0.8mM,
P=588mM
800
Fe(III)=81mM,
Cu(II)=16mM,
P=588mM
700
Fe(III)=81mM,
Cu(II)=0.8mM,
P=588mM
600
500
0
1
2
3
Time [h]
4
5
6
Figure 16. Example for the results obtained in the catalytic oxidation experiments.
Concentration of Fe(II) as a function of time in four different experiments. pH1.0
64
In most of the studies that addressed the spontaneous oxidation rate of Fe(II) by oxygen
at low pH levels (pH < 2) the kinetic equation was found to be first-order with respect to
oxygen, first or second-order with respect to Fe(II) and independent of pH.
For the case that the rate is first order in Fe(II) the most widely accepted mechanism
is the one presented by King and Davidson (1958):
a ,b
Fe( II ) + O2 + H + ←→
Fe( III ) + HO2
•
•
(46a,b)
Fe( II ) + HO2 + H + → Fe( III ) + H 2 O2
(47)
2 Fe( II ) + H 2 O2 + 2 H + → Fe( III ) + 2 H 2 O
(48)
Equation (46a) is the rate determining step and Eq. (48) proceeds very rapidly.
The reaction order with respect to Fe(II) concentration in the catalytic oxidation
experiments was assessed under the assumption of zero-, first- and second-order kinetics.
The best fit was used to calculate the oxidation rate according to 40% reaction
completion, i.e. the time needed to oxidize 400 mg/l of Fe(II) (the initial concentration in
all the experiments was around 1000 mg/l). Reducing the Fe(II) concentration below 400
mg/l has no advantage from the standpoint of the LRSR process. Second-order kinetics
was found to be the most suitable for 40 out of 60 experiments. In 15 out of 60
experiments a zero-order dependency was the most fitting. No clear correlation could be
observed between the reaction order (with respect to the Fe(II) concentration) and other
components in the aqueous phase (i.e., iron, phosphate and copper concentrations).
The effects of the following factors on the oxidation rate were investigated in the
catalytic Fe(II) oxidation experiments: pH, initial Fe(III) concentration, total dissolved
phosphate concentration (PT) and initial Cu(II) concentration. Since the counter anion
of all the reagents was SO42- (and also H2SO4 was used to adjust the pH of the tested
solutions) the effect of total sulfate concentration (ST) was also examined.
4.2.6.1. Effect of pH
The pH was measured at the beginning and end of each experiment. The average
change in pH was ±0.03 and maximal change was ±0.10. The pH was expected to
increase with time, since Fe(II) oxidation consumes alkalinity (Eq. (19)). However,
aqueous solution at such low pH levels has a high buffering capacity. Table 10 lists
the conditions of experiments that were conducted at pH different then 1.0: the first
four experiments were conducted at pH~1.6 and the rest – at pH~0.8. For each of
65
these experiments, a parallel experiment at pH1.0 was conducted (i.e. with the same
initial concentrations of Fe(III), H3PO4 and Cu(II)).
Table 10. The conditions of catalytic oxidation experiments that were conducted at
pH values other than 1.0 and representative Fe(II) oxidation rates
Concentration [mM]
pH
Representative Fe(II) oxidation rate
[mg-Fe/l/h]
PT
ST
Cu(II)
Fe(III)
A
613
31
7.9
0
1.60
684
B
807
31
7.9
0
1.55
888
C
807
62
0.0
0
1.59
57
D
613
94
0.0
0
1.64
28
E
743
281
0.8
161
0.78
25
F
291
312
0.8
161
0.83
9
G
452
312
0.8
161
0.80
14
H
581
343
0.8
161
0.75
18
I
0
499
0.8
161
0.77
6
J
387
530
7.9
322
0.78
4
K
387
530
0.0
322
0.77
5
L
613
530
7.9
322
0.72
10
M
807
530
7.9
322
0.67
16
A comparison of the representative ferrous oxidation rates (in the specified Fe(II)
concentration interval) at different pH values is presented in Fig. 17 and Fig. 18. It is
clear from Fig. 17 that at pH~1.6 the rate of Fe(II) oxidation is one order of
magnitude higher than at pH 1.0. Although a higher oxidation rate is ostensibly
advantageous, precipitation of Fe(III) oxides at pH~1.6 is rapid, making this pH
infeasible for long term operation of the LRSR process. An attempt to conduct
experiments similar to experiments A to D in the presence of Fe(III) (161 mM) failed,
because a massive Fe(III) precipitation occurred (the solution was clear only when pH
was lowered below pH1.3).
66
1000
684
888
pH>1.0
-d[Fe(II)]/dt [mg Fe/l/h]
242
pH=1.0
78
100
57
28
7
10
6
1
A
B
C
Experiments
D
Figure 17. Comparison between rates of Fe(II) oxidation at pH~1.6 and pH~1.0 under
various conditions (see Table 10 for specific operational conditions). y-axis is in
logarithmic scale
Since precipitation could not be avoided also at pH1.0 (see 4.1.2) in the presence of
phosphate the oxidation rate was investigated also at lower pH (pH0.8, experiments E
to M in Table 10). The effect of lowering the pH below 1.0 on the oxidation rate was
not conclusive, as shown in Fig. 18.
-d[Fe(II)]/dt [mg Fe/l/h]
35
pH=1.0
30
pH<1.0
25
20
15
10
5
0
E
F
G
H
I
J
Experiments
K
L
M
Figure 18. Comparison between rates of Fe(II) oxidation at pH~0.8 and pH~1.0 under
various conditions (see Table 10 for specific operational conditions)
Although the oxidation rates at pH0.8 were not much lower than the rates at pH1.0 it
was decided to set the working pH at pH1.0. This was done for two reasons: first, the
67
precipitation was minimized but not avoided even at pH0.8, and second, the lower pH
necessitates the use of even more resistant equipment, which makes the process less
cost effective.
All experiments presented from here onward were conducted at pH1.0 (±0.1). As
mentioned earlier there was no apparent change in pH during the experiments.
4.2.6.2. Effect of initial concentration of Fe(III)
An investigation of the H2S(g) reactive absorption efficiency in a bubble column
reactor at low pH was reported only in one study, by Gendel (2007), who applied the
following conditions: pH 1.72 and 81 or 161 mM Fe(III). Gendel (2007) found, like
others before him (see Subsection 1.2.4.2), that an increase in the concentration of
Fe(III) results in an increase in the reactive-absorption efficiency. Efficiencies higher
than 95% and 75% were attained using Fe(III) concentrations of 161 and 81 mM
respectively and air flux of 1 m3/m3/min. Since a bubble column reactor was used in
the current study, the Fe(III) concentrations used by Gendel (2007) were adopted for
the examination of their effect on the Fe(II) oxidation rate.
Experiments with similar initial concentrations of Cu(II) and PT, were gathered into a
series. In Fig. 19 the oxidation rates of 6 of these series are presented. The initial
[Fe(III)]:[Fe(II)] ratio were 0, 4.5, 9 and 18. Error bars were added in order to
emphasize the change in Fe(III) concentration during the experiments as a result of
Fe(II) oxidation.
It is clear from Fig. 19 that Fe(III) has an inhibiting effect on the Fe(II) oxidation rate
(exponential tendency seems to explain well the retarding effect of Fe(III)
concentration, 0.636<R2<0.998). From a chemical standpoint the retarding effect of
Fe(III) may be explained by Eq. (46b), i.e. reduction of Fe(III) by HO2●. Pham and
Waite (2008), who investigated the Fe(II) oxidation by O2 in natural waters,
concluded that the reduction of Fe(III) is important particularly under conditions
where precipitation of Fe(III) is minimal. This seems to explain well the decrease in
Fe(II) oxidation rate that was obtained with an increase in the initial [Fe(III)]/[Fe(II)]
ratio, because the experiments were conducted at pH1.0, where precipitation of Fe(III)
is relatively low.
68
-d[Fe(II)]/dt [mg Fe/l/h]
250
[Cu]=0.8 mM
[Cu]=1.6 mM
[Cu]=3.1 mM
[Cu]=6.3 mM
[Cu]=7.9 mM
[Cu]=15.7 mM
200
150
PT = 581 mM
100
50
0
0
50
100
150
200
250
300
350
Fe(III) initial concentration [mM]
Figure 19. Fe(II) oxidation rate as a function of the initial Fe(III) concentration with
various concentrations of Cu(II). Error bars present the change in Fe(III)
concentration during the experiments
Although the oxidation rates were relatively high when no Fe(III) was initially
present, these results are meaningless with regard to the H2S(g) removal process since
a high concentration of Fe(III) is essential for efficient reactive-absorption of H2S(g).
From the standpoint of the LRSR process it might be more practical to use an Fe(III)
concentration of 81 mM (4.5 g Fe/l) in the working solution rather than 161 mM (9
g/l), since the Fe(II) oxidation rate in the presence of 81 mM Fe(III) is almost twice as
high as the rate in the presence of 161 mM and the H2S(g) reactive absorption
efficiency can still be quite high, according to the results of Gendel (2007).
The synergistic effect of Fe(III) and Cu(II) is discussed in the following Subsection.
4.2.6.3. Effect of initial concentration of Cu(II)
As mentioned in Section 1.3.3 the catalytic effect of Cu(II) on Fe(II) oxidation is well
known. Cher and Davidson (1955) observed that the catalytic effect in phosphoric
acid solutions was significant for Cu2+ concentration up to 10−5 M, but there was little
change with increasing the Cu2+ concentration above 10−3 M. This limited catalysis by
Cu2+ was also observed by Zhang et al. (2000), who showed that the rate of Fe2+
69
oxidation is significantly enhanced by adding 0.02 M CuSO4, but a further increase in
the Cu2+ concentration to 0.1 M results in no additional change (Fig. 20).
Figure 20. Effect of Cu2+ on Fe2+ oxidation by O2 (0.1 M FeSO4; pH=1.5; 80°C).
(Source: Zhang et al., 2000)
Similar results were obtained in the current work. In Fig. 21 the oxidation rates of
three oxidation experiment series (each series with a different initial concentrations of
Fe(III)) are presented. It is noticeable that the Fe(II) oxidation rate increased with an
increase in Cu(II) concentration, but above ~4mM of Cu(II) and in the presence of
Fe(III) the accelerating effect leveled off.
This catalytic action is believed to be initiated by the reaction given in (49a) (Cher
and Davidson, 1955; Zhang et al., 2000):
a ,b
Fe 2+ + Cu 2+ ←→
Fe 3+ + Cu +
(49a,b)
and followed by the reactions:
Cu + + O2 → CuO2
+
+
CuO 2 + H + → Cu 2+ + HO2
(50)
•
(51)
It was assumed that the first reaction (Eq. (49a)) is the rate determining step and that
HO2• further reacts with Fe(II) as in the Weiss mechanism (Huffman and Davidson,
1956). It was concluded that the catalytic action of copper ion on the oxidation of Fe2+ by
O2 relies on the fact that Cu+ reacts with O2 more readily than Fe2+ to generate HO2• and
H2O2, and that the oxidation of Fe2+ ion by H2O2 is faster than by O2.
70
300
-d[Fe(II)]/dt [mg Fe/l/h]
[Fe(III)]=0
250
a
[Fe(III)]=81 mM
[Fe(III)]=161 mM
200
150
100
50
0
0
2
4
6
8
10
12
14
16
Cu(II) concentration [mM]
-d[Fe(II)]/dt [mg Fe/l/h]
80
b
60
40
[Fe(III)]=0
20
[Fe(III)]=81 mM
[Fe(III)]=161 mM
0
0
2
4
6
8
10
12
Cu(II) concentration [mM]
14
16
Figure 21. Fe(II) oxidation rate as a function of the initial Cu(II) concentration with
various concentrations of Fe(III). The × signs represent oxidation experiments that
were conducted with pure oxygen (as opposed to oxidation with air). PT=581mM
However, according to Zhang et al. (2000) Cu(I) further reacts with HO2• to form
H2O2 (Eq. (52a)):
•
a ,b
Cu + + HO2 + H + ←→
Cu 2+ + H 2 O2
(52)
This phenomenon is apparently due to competing oxidation of Cu2+ according to
Equations (52b) and (53).
•
HO 2 + Cu 2+ → O2 + Cu + + H +
71
(53)
Zhang et al. (2000) also found that a high Cu2+ concentration (five times higher than the
initial Fe2+ concentration) does accelerate the initial rate of Fe2+ oxidation, but has little
effect once significant Fe3+ has been produced in the reaction mixture. The explanation
for this is as follows: it appears that at relatively high initial concentrations of both Cu2+
and Fe2+, the oxidation of Fe2+ occurs via the generation of Cu+ (Eq. (49a)), and the
subsequent production of H2O2. However, as the oxidation proceeds, Fe3+ builds up and
the reverse reaction between Fe3+ and Cu+ (Eq. (49b)) becomes predominant. Thus, Fe3+
competes with O2 to oxidize Cu+ and as a result, lessens or disables the formation of
CuO2+ that is the key intermediate of the proposed mechanism.
In order to evaluate the effect of O2 on Cu(II) catalysis, three experiments were
conducted with pure oxygen. The solutions contained 0.8 mM Cu(II) and the initial
Fe(III) concentrations were 0, 0.08 and 0.16 M (represented as × signs in Fig 21). In the
absence of Fe(III) the oxidation rate was similar to the oxidation rate attained with air
(see Fig. 21b), which means that the oxidation rate is not limited by the concentration of
O2. However, in the presence of Fe(III) the oxidation rate was higher than the oxidation
rate with air (Fig. 21b). This result seems to indicate that the increase in O2 concentration
increased the oxidation rate of Cu+ by O2 rather then the competing reaction with Fe3+,
which resulted in a faster total-oxidation rate of Fe(II).
Cher and Davidson (1955) developed the following rate equation according to their
findings in phosphoric acid solution:
[
][ ]
[ ]
PO2 
Fe 2+ Cu 2+ PO2

k1 Fe 2+ PO2 1 + k1
+
k
 3
Fe 3+ 
Fe 3+
d [Fe( II )]

−
=
PO2
dt
Cu 2+
1 + k2
+
k
4
Fe 3+
Fe 2+
[
]
[
]
[
]
[
[
]
]
(54)
The saturation effect with respect to Cu2+ and the retarding effect of Fe3+ were observed
also in sulfuric acid solutions (Lowson, 1982). However, the rate constant for the first
step in the cupric ion catalyzed reaction (Eq. (49a)) in phosphoric acid is two-orders of
magnitude higher than in sulfuric acid solution (Huffman and Davidson, 1956). The
catalytic effect of Cu2+ in acidic ferrous chloride solutions had also been noted (Lowson,
1982).
To summarize, the catalytic effect of Cu(II) is limited in the presence of Fe(III), and in
the range of conditions that was examined in the current study, increasing the Cu(II)
concentration above 4 mM results in an insignificant increase in the rate of Fe(II)
72
oxidation. The use of pure oxygen acts to increase the Fe(II) oxidation rate, however it is
not cost effective and cannot be considered a process option..
4.2.6.4. Effect of total phosphate concentration
As discussed in Section 1.3.3 phosphate was found to increase the Fe(II) oxidation
rate due to its complexing effect with the Fe(III) product (Cher and Davidson, 1955;
King and Davidson, 1958).
In Fig. 22 the oxidation rates of seven oxidation experiment series (each series with
different initial concentrations of Fe(III) and Cu(II)) are presented. It is clear from the
results that phosphate has a catalytic effect on the Fe(II) oxidation rate. A second
order dependency on the total phosphate concentration (PT) was found to fit well the
results, irrespective of the Fe(III) and Cu(II) concentrations (0.926<R2<1.000).
In order to better understand the results, an analysis of the aqueous phase distribution
of phosphate species in solution with respect to PT was made with the software
MINEQL+. The sets of the input data for this theoretical analysis are given in Table
11. The distribution of the Fe(III) species and Fe(II) species for set D as a function of
the PT are presented in Figures 23 and 24 respectively.
Table 11. Composition of solutions that were analyzed by the MINEQL+ software for
the distribution of phosphate species
Concentrations [mM]
Fe(II)
Fe(III)
Cu(II)
H2SO4*
Set A
18
0
0
100 to 35
Set B
18
0
8
75 to 45
Set C
18
161
0
400 to 200
Set D
18
161
8
400 to 300
* At a higher H3PO4 concentration in solution less H2SO4 was needed to reach pH1.0
73
80
[Fe(III)]=0, [Cu]=0
[Fe(III)]=0, [Cu]=7.9
[Fe(III)]=81, [Cu]=0.8
[Fe(III)]=81, [Cu]=1.6
[Fe(III)]=81, [Cu]=7.9
[Fe(III)]=161, [Cu]=0.8
[Fe(III)]=322, [Cu]=7.9
-d[Fe(II)]/dt [mg Fe/l/h]
70
60
50
in mM
40
30
20
10
0
0
100
200
300
400
500
600
700
Total phosphate concentration [mM]
800
900
Figure 22. Fe(II) oxidation rate as a function of PT with various concentrations of
Fe(III) and Cu(II)
Figure 23. Distribution of the Fe(III) species as a function of PT, in the presence of 1
g Fe(II)/l (18 mM), 9 g Fe(III)/l (161 mM) and 0.5 g Cu(II)/l (8 mM). pH1.0. Semilogarithmic scale
74
Figure 24. Distribution of the Fe(II) species as a function of PT, in the presence of 1 g
Fe(II)/l (18 mM), 9 g Fe(III)/l (161 mM) and 0.5 g Cu(II)/l (8 mM). pH1.0. Semilogarithmic scale
The analysis show that the dominant species of Fe(III) in the presence of phosphate is
FeH2PO42+ (at PT > 0.12 M), and the dominant species of Fe(II) is FeH2PO4+ (at PT > 0.18
M). In addition second-order dependence was found to fit well the increase in
concentrations of the species FeH2PO4+ and the ion H2PO4- with respect to PT (R2=0.9986
and 0.9982 respectively). The tendency was similar in the other simulations as well (sets
A to C).
Cher and Davidson (1955) and King and Davidson (1958) demonstrated that the
oxidation of ferrous sulfate by molecular oxygen in phosphoric acid and pyrophosphoric
acid solutions (at pH ~1-2, with the ionic strength adjusted to 1.0-1.1 with sodium
perchlorate, and at 30°C) was second order with respect to H2PO4- and first order with
respect to H2P2O72-, yielding the following overall rate law
−
d [Fe( II )]
−
= k1 Fe 2+ PO2 H 2 PO4
dt
[
] [
] + k [Fe ]P [H P O ]
2
2−
2+
2
O2
2
2
7
(55)
King and Davidson (1958) concluded that the second-order dependence on phosphate
was not due to the equilibrium
−
2 H 2 PO4 ↔ H 2 P2 O7
2−
+ H 2O
(56)
75
Thus, the second order dependency with respect to H2PO4- concentration, which was
found in the current work, is corroborated by literature.
No inhibition by Fe(III) was observed by Cher and Davidson (1955), and it was
suggested that the complexing effect of Fe(III) by phosphate decreased the rate of
Fe(III) reduction by HO2• (Eq. 46b). However, there was no Fe(III) at the beginning
of the experiment of Cher and Davidson (1955), while in current work the initial
concentration of Fe(III) was high and as a result inhibition was observed right from
the start of the experiments.
All in all, the results show a positive effect of phosphate on the Fe(II) oxidation rate,
probably related to the concentration of the H2PO4- ion. The drawback of using a high
PT is its rapid precipitation with Fe(III) even at the low pH value practiced in the
current work (see Section 4.1.2.).
4.2.6.5. Effect of total sulfate concentration
The oxidation of ferrous sulfate in sulfuric acid solution has been described as follows
(Chmielewski and Charewicz, 1984; Rönnholm et al., 1999):
4 FeSO4 + O2 + 2 HSO4 + 2 H 3 O + → 2 Fe2 (SO4 )3 + 4 H 2 O
−
(57)
Most of the previous studies indicate that at ambient temperatures and in sulfuric acid
solutions below pH2, the rate of reaction is very slow, independent of pH, first order
with respect to the partial pressure of oxygen and second order with respect to ferrous.
Combined, these results are summarized in the following general rate equation
(Lowson, 1982; Chmielewski and Charewicz, 1984):
−
2
d [Fe( II )]
= k Fe 2+ PO2
dt
[
]
(58)
Huffman and Davidson (1956) observed an increase in the reaction rate with sulfate
concentration, at a fixed pH and ionic strength. .The rate was found to be independent
of dissolved ferric ion provided that additional sulfuric acid is added; otherwise the
product slows down the reaction rate (Macejevskis and Liepina, 1965, quoted by
Lowson 1982). The presence of sulfate ions at near-neutral pH (as discussed in 1.3.2)
decreased the oxidation rate, presumably because of the formation of the FeSO40
complex, which does not oxidize rapidly as compared to the complexes with
hydroxyls and carbonates. It is possible that at low pH levels, where the
concentrations of hydroxyls and carbonates are negligible, the oxidation of FeSO40
76
becomes significant. This hypothesis is supported by Dreisinger and Peters (1989)
who reported on higher reactivity of the ferrous sulfate ion pair relative to the free
ferrous ion, and Willix, (1963) and Ciavatta et al., (2002), who found that at low pH
the product FeSO4+ is stable. Huffman and Davidson (1956) reported that the
oxidation is accomplished by both catalyzed and un-catalyzed reactions, as described
in the following rate expression:
−
2
d (Fe( II ) )
2
= k u Fe 2+ PO2 + k s (FeSO4 ) Fe 2+ PO2 = k t (Fe( II ) ) PO2
dt
(
)
(
)
(59)
In the current study the total sulfate concentration (ST) in the tested solutions was
between 0.04 and 0.53 M (1.1 to 16.9 g S/l), resulting from the concentrations of the
reagents added ((NH4)2Fe(SO4)2•6H2O, CuSO4, Fe2(SO4)3 and H2SO4). Fe(II) oxidation
rates as a function of ST are presented in Fig. 25. The results are quite scattered, thus it
can be only concluded that sulfate has a minor effect on the oxidation rate of Fe(II). The
somewhat decreased reaction rate observed with an increase in ST can be attributed to the
increase in Fe(III) concentration, that originates from the reagent Fe2(SO4)3.
300
in mM
P=0, Cu=0
P=0, Cu=7.9
P=452, Fe(III)=81
P=581, Fe(III)=0
P=581, Fe(III)=81
P=581, Fe(III)=161
-d[Fe(II)]/dt [mg Fe/l/h]
250
200
150
100
50
0
0
50
100
150
200
250
300
350
400
450
Total sulfate concentration [mM]
Figure 25. Fe(II) oxidation rate as a function of ST with various concentrations Fe(III),
Cu(II) and PT
77
Summary: In the investigation of the catalytic oxidation of Fe(II) by oxygen the
following results were obtained:
The Fe(II) oxidation rate decreases with a decrease in Fe(III) concentration. This
indicates that a low Fe(III) concentration is requested in order to enhance Fe(II)
oxidation kinetics. However, a high concentration of Fe(III) is needed for efficient
reactive-absorption of H2S(g), and thus this option is not practical.
The Fe(II) oxidation rate is accelerated in the presence of Cu(II). The use of pure
oxygen rather than air may further increase the oxidation rate. However, H2S reacts
with Cu(II) and precipitate as CuS. Under the investigated condition CuS did not
oxidize to S0 by Fe(III) as reported by Zhang and Tong (2006).
Moreover, the Fe(II) oxidation rate is accelerated in the presence of phosphate. Yet,
the phosphate concentration that is needed in order to achieve a reasonable oxidation
rate (higher or equal to the oxidation rate of H2S by Fe(III)) is not feasible, since
massive precipitation of P-Fe(III) species will occur.
At pH1.0, in the presence of 4.5-9.0 g Fe(III)/l (81-161 mM) and less than 14 g P/l
(452 mM) the maximal oxidation rate achieved was 7.5 mg Fe(II)/h. With this rate it
is possible to treat up to 2.3 mg H2S/h, as calculated according to Eq. (60).
rH 2 S( g )
 g H 2S 
 g Fe ( II ) 
rFe ( II ) 

 × M .W . H 2 S 
 g H 2 S  (60)
 g Fe ( II ) 
 gmol H 2 S 
 h 

=
= rFe ( II ) 

 × 0.30515
 g Fe ( II )   gmol Fe ( II ) 
 h 
 g Fe ( II ) 
M .W . Fe ( II ) 
×E
 × 2
 gmol Fe ( II )   gmol H 2 S 
where rH 2 S( g ) is the rate of H2S(aq) oxidation, rFe ( II ) is the rate of Fe(II) oxidation and E
is the reactive-absorption efficiency of H2S(g) (assumed to be 1).
All in all, the oxidation of Fe(II) by O2 at pH1.0 can be catalyzed by phosphoric acid
and copper ions. However, this procedure cannot be considered feasible for the
process of H2S(g) removal.
78
4.3. Electrochemical Fe(II) oxidation as part of the LRSR process
4.3.1. H2S(g) reactive-absorption efficiency
A PVC pipe was used as a bubble column reactor (internal diameter 158 mm). A
sintered-glass plate diffuser (5 inches diameter) with a pore size distribution of 150 200 µm was used to generate the bubbles. The effect of pH and chloride concentration
was tested with an air flux of 1.1 m3air/min/m3solution and the results are presented in
Fig. 26 and Fig. 27 as a function of the inlet H2S(g) concentration. It seems that both
increase in pH and increase in the chloride concentration in solution have a positive
effect on the efficiency of H2S(g) reactive-absorption. However, the effect of chloride
was much stronger in the range of conditions tested: while increasing the pH from 1.0
to 1.7 elevated the reactive-absorption efficiency from 80% to 89% in average, the
addition of 0.28 M Cl- elevated the efficiency from 18% to 70% in average.
H2S(g) reactive-absorption efficiency
[%]
95
pH=1.74
pH=1.38
pH=1.00
90
85
80
75
0
10
20
30
40
H2S(g) concentration [ppm]
50
60
Figure 26. Effect of pH on the H2S(g) reactive-absorption efficiency. [Cl-] = 0.85M,
air flux = 1.1 m3/m3/min
79
H2S(g) reactive-absorption efficiency
[%]
90
80
70
0 M Cl0.28 M Cl0.85 M Cl1.41 M Cl-
60
50
40
30
20
10
0
0
10
20
30
40
H2S(g) concentraion [ppm]
50
60
Figure 27. Effect of Cl- concentration on the H2S(g) reactive-absorption efficiency.
pH1.0, air flux = 1.1 m3/m3/min
Fig. 28 presents the distribution of the Fe(III) species as a function of pH in the
presence of 0.85 M Cl- (30 g Cl/l) (analysis using MINEQL+). The concentrations of
the species FeOH2+, Fe(OH)2+ and Fe2(OH)24+ increase with an increase in pH. The
concentration of the species Fe3(OH)45+ also increases significantly, from 10-9.1 to 106.4
M. The same tendency was found also in solution in the absence of Cl-. The species
FeOH2+ was reported by Asai et al. (1990) and Ebrahimi et al. (2003) to oxidize
H2S(aq). Ebrahimi et al. (2003) hypothesized that the species Fe3(OH)45+ and
Fe2(OH)24+ also oxidize H2S(aq) (see Subsection 1.2.4.2). Thus it can be safely
concluded that the increase in the reactive-absorption efficiency with pH is a result of
the increase in the concentrations of the Fe(III) hydroxyl complexes.
The effect of chloride concentration on the H2S(g) reactive-absorption efficiency was
also analyzed by MINEQL+ while pH was maintained constant. The results are
presented in Fig. 29.
80
Figure 28. Distribution of Fe(III) species as a function of pH in the presence of Cl-
(0.85 M). Semi-logarithmic scale.
Figure 29a. Distribution of Fe(III) species as a function of chloride concentration.
Semi-logarithmic scale. pH1.0
81
H2S(g) reactive-absorption efficiency
[%]
Figure 29b. Distribution of Fe(III) species as a function of chloride concentration.
Normal scale. pH1.0
According to Fig. 29a the concentrations of the hydroxyl complexes, which were
found to oxidize H2S(aq), decrease with the increase in total-Cl- concentration. It could
have been assumed that the H2S(g) reactive-absorption efficiency would decrease with
the increase in Cl- concentration. However, the reactive-absorption efficiency of
H2S(g) in the presence of Cl- was found to be much higher than the efficiency in the
absence of Cl- (Fig.s 27 and 29b). Even at low total-[Cl-], the concentrations of the
Fe3+Cl- complexes are high and increase with an increase in total-[Cl-]. It is thus
concluded that the Fe3+Cl- complexes can also oxidize well H2S(aq). No information
can be found in the literature on reactive-absorption of H2S(g) in the presence of
chlorides, and the observation that the Fe3+Cl- complexes can oxidize H2S(aq) is new.
In Fig. 29b the average values of H2S(g) reactive-absorption efficiencies, for each one
of the experiments presented in Fig. 27, are also presented. It is noticeable that the
dependency of the reactive-absorption efficiency on the total-[Cl-] is similar to the
dependency of the concentration of the FeCl2+ species on total-[Cl-]. It is thus not
unlikely that the dominant oxidizing complex of H2S(g) at pH1.0 is FeCl2+.
The main conclusion from this section is that the use of a solution that does not
contain chlorides is not applicable for the LRSR process at pH1.0. A second
82
conclusion is that increasing the Cl- concentration beyond 30 g/l (0.85 M) does not
improve the reactive-absorption efficiency of H2S(aq) and seems thus to be
superfluous.
4.3.2. Practical precipitation potential in the working solution
The effect of pH and chloride concentration was tested. The compositions of the
tested solutions are listed in Table 12.
Table 12. Conditions of solutions for precipitation experiments (electrochemical
oxidation) and mass of precipitates after 152 days
Bottles
No.
Concentrations [mM]
Fe(II)
A
-
Fe(III)
Cl
161
846
(=9.0 g Fe/l)
(=30 g Fe/l)
161
pH
Mass of precipitate
[mg per l-solution]
1.0
8
846
1.4
56
161
846
1.7
528
161
282
1.0
31
E
161
282
1.4
34
F
81
846
1.0
45
B
C
D
18
(=1 g Fe/l)
From a thermodynamic standpoint and based on MINEQL+ simulations all the
considered solutions were projected to precipitate Fe(III) solids to a degree that the
soluble FeT concentration remaining in solution at equilibrium would be very low. For
example, in all the experiments which were conducted at pH1.0 (A, D and F) more
than 94% of the Fe(III) was supposed to precipitate as part of the solid Fe(OH)2.7Cl0.3.
In practice, however, the pH value was found the most dominant parameter in the
formation of ferric precipitates under the investigated conditions.
After a period of 51 days all solutions seemed visually clear. However, on the walls of
bottle C (0.85 M Cl- at pH1.7) a white precipitate could be observed (Fig. 30). No
change in pH was detected. After 138 days from the beginning of the experiment,
precipitation has increased in bottle C and small suspended clusters could be seen in
solution (although no turbidity was observed). In bottle B (0.85 M Cl- at pH1.4) very
83
slight precipitation was observed. After 152 days the solutions (<500 ml) were filtered
through a 0.45 µm filter paper and the filters were dried in 105°C for 24 hours. The
mass of precipitates that were found are presented in Table 12. After 152 days a
significant amount of precipitates (528 mg per l-solution) were observed in bottle C.
In the bottles that were held at pH1.0 and pH1.4 the amount of precipitates was much
lower (more than one order of magnitude). Yet, at pH1.4 the amount of precipitates
was higher than at pH pH1.0.
It can be concluded that although precipitation was anticipated by thermodynamic
data, the kinetics of precipitation was low and affected by pH. The pH of solution
should not be higher than pH1.4 for a long-term LRSR process. Working at pH1.0
provides a safety factor.
C
B
Figure 30. Bottles from precipitation experiments after 51 days. B (pH1.4) and C
(pH1.7). [Cl-]=0.85 M
4.3.3. Fe(II) oxidation rate in the electrochemical oxidation process
The majority of the experiments were performed using a constant current (±0.1 amp)
and the voltage values were measured. Otherwise, voltage was kept constant and the
current was measured. All the electro-oxidation experiments were conducted at
pH1.0.
In earlier experiments (Subsections 4.2.3.1 and 4.2.3.2) the cathode was made of
titanium coated with ruthenium oxide (Ti/RuO2). Later on it was decided to replace
the electrode to bare Ti, since the RuO2 coating started to detach from the electrode
after a few hours of operation. A bare Ti electrode was used in the experiments shown
in Subsections 4.2.3.3 and 4.2.3.4.
84
4.3.3.1. Direct electrooxidation
The rate of direct oxidation was evaluated and compared to the rate of indirect
oxidation. The conditions of the experiments are listed in Table 13 and the results are
shown in Fig. 31. The total iron concentration in the experiments was 0.179 M (10 g
Fe/l). Although the Fe(III) concentration in the proposed process are designed to vary
inside the 9-8 g Fe/l interval, initial Fe(II) and Fe(III) concentrations were both set at
5 g Fe/l (0.090 M) in order to obtain the influence of the Fe(III) concentration on the
Fe(II) electro-oxidation rate.
It is noticeable from Fig. 31 that the oxidation rate of Fe(II) is higher at lower
[Fe(III)]:[Fe(II)] ratio. In fact, there were two tries to perform direct oxidation in
solution with an initial [Fe(III)]:[Fe(II)] ratio of 9:1. These experiments were stopped
after a short period (about 30 minutes) since the concentration of Fe(II) increased
rather than decreased with time. It was concluded that Fe(III), being at such high
concentration, competes with H+ on reduction on the cathode surface and thus under
such conditions the back reduction reaction of Fe(III) to Fe(II) was faster than the
forward reaction (Fe(II) → Fe(III)). Yet, high ratio of [Fe(III)]:[Fe(II)] is essential in
order to achieve efficient reactive-absorption of H2S(g) (see 1.2.4.2).
The Fe(II) oxidation rate recorded for the direct oxidation was significantly lower
than the rates attained in the indirect oxidation experiments: 0.22 g/h in first two hours
for direct oxidation compared to 1.41 g/h for indirect oxidation. Nevertheless, these
rates are much higher than the rates that were obtained in the catalytic oxidation
(Section 4.1.3).
Since the preliminary experiments of indirect oxidation showed higher oxidation rates
than direct oxidation and since the reactive-absorption efficiency of H2S(g) was very
low in the absence of chlorides (Section 4.2.1), the indirect oxidation option seemed,
from the inception, much more promising as the regenerative step for Fe(III) within
the LRSR process at pH1.0.
4.3.3.2. Indirect electrooxidation - effect of anode to cathode surface area ratio
(Sa:Sc)
High anode to cathode surface area ratio (Sa:Sc) may suppress the Fe(III) back
reduction on the cathode, which is important because of the need for high Fe(III)
85
concentration in the LRSR solution. On the other hand, high Sa:Sc may suppress the
wanted H2 evolution on the cathode.
Two surface area ratios (Sa:Sc) were examined: 7.1 (smaller cathode) and 1.7 (bigger
cathode). The conditions of these experiments are listed in Table 13 and the results are
shown in Fig. 31.
It is very clear from Fig. 31 that the net Fe(II) oxidation rate was higher in the
experiments with lower Sa:Sc ratio (bigger cathode). This means that H2 evolution was
more dominant than Fe(III) back reduction on the cathode surface, and increasing the
cathode surface allowed higher H2 evolution rate and consequently higher Fe(II)
oxidation rate. This is supported by Alvarez-Gallegos and Pletcher (1998) and Bisang
(2000), who reported that the main reaction on the cathode around pH2 was H2
evolution (see Subsection 1.4.2.1).
Table 13. Conditions of electrochemical oxidation experiments conducted with a
Ti/RuO2 anode
Direct /
indirect
oxidation
Direct
Indirect
Indirect
Indirect
Indirect
Initial
Current
Voltage
pH
[amp]
[V]
1.01
0.3
2a
0.95
1.03
Cl-
Cathode
concentration
dimensions
[M]
[cm2] b
0 : 10
0
0.72 × 9.15
Fe(III):Fe(II)
a
1
3.5
5:5
0.85
0.72 × 9.15
1
3
a
5:5
0.85
0.72 × 9.15
3
a
5:5
0.85
3.0 × 9.14
2.6
5:5
0.85
3.0 × 9.14
1.0
2.2
1.0
a
2
a
was kept constant during the experiment
b
anode dimensions 5.1 x 9.15 cm2 in all experiments
86
Fe(II) concentration [g Fe/l]
Figure 31. Concentration of Fe(II) as a function of time in the experiments conducted
with Ti/RuO2 cathode
4.3.3.3. Indirect electrooxidation - effect of current density
The range of currents that were investigated in literature is between 0.2 and 4.2 kA/m2
(Robertson et al., 1983; Kelsall, 1984; Rudolf et al., 1995; Khelifa et al., 2004). An
increase in the active chlorine production rate as a result of increase in the current
density was reported by Rudolf et al. (1995), Kraft et al., (1999) and Khelifa et al.
(2004). The effect of current density in the current work was tested with a Ti cathode
and Ti/RuO2 anode, both with dimensions of 9.15×5.1 cm2. The range of current
densities was between 0.43 and 1.71 kA/m2 (2 to 8 A with 9.15×5.1 cm2 electrodes).
The chloride concentration in all the experiments was 0.85 M. This concentration was
chosen according to the results of the H2S(g) reactive-absorption efficiency
experiments (Section 4.2.1).
The interesting interval of Fe(II) concentration is between 2 to 1 g Fe/l from the
standpoint of the proposed LRSR process (and Fe(III) concentration is 8-9 g/l). Thus,
the Fe(II) oxidation rate was calculated in this interval by a linear regression. For
example, the incline of the linear regression for the experiment conducted with 0.43
kA/m2 was -1.1356 g Fe(II)/l/h (Fig. 32). The solution volume was 1.5 l, thus the
Fe(II) oxidation rate was 1.70 g Fe(II)/h, as seen in Fig. 33.
87
5
0.43 kA/m2
0.54
Fe(II) concentration [g Fe/l]
4
kA/m2
0.64 kA/m2
0.86 kA/m2
1.07 kA/m2
1.29 kA/m2
1.71 kA/m2
3
2
y = -1.1356x + 3.9024
R2 = 0.9946
1
0
0.0
0.5
1.0
1.5
2.0
Time [h]
2.5
3.0
3.5
Figure 32. Concentration of Fe(II) as a function of time in the indirect
72
7
64
6
56
48
40
4
32
3
24
2
-d[Fe(II)]/dt
Voltage
1
0
0.0
16
Current efficiency [%]
8
5
Voltage [V]
-d[Fe(II)]/dt [g Fe/h]
electrochemical experiments at different current densities. [Cl-]=0.85 M
Current efficiency
0.5
1.0
1.5
2
Current density [kA/m ]
8
0
2.0
Figure 33. Effect of current density on Fe(II) oxidation rate, voltage and current
efficiency. [Cl-]=0.85 M
88
Fig. 33 shows the Fe(II) oxidation rates (in the interesting interval) and the current
efficiencies as a function of the current densities. Higher Fe(II) oxidation rates were
achieved for higher current densities. The highest Fe(II) oxidation rate recorded was
7.1 gFe/h, and it corresponded to current efficiency of 42.8%. The highest current
efficiency (59%) was achieved at 0.64 kA/m2.
4.3.3.4. Indirect electrooxidation - effect of chloride concentration
Khelifa et al. (2004) examined the effect of chlorine concentration in the range of 6 to
150 g-Cl/l and reported on an increase in active chlorine production with an increase
in chloride concentration. The same tendency was observed by Krstajić et al. (1991)
who examined the hypochlorite production rate for NaCl solutions with
concentrations of 14.6 g/l, 19.28 g/l and 29.21 g/l. However, it was also reported by
Khelifa et al. (2004) that at NaCl concentration above 100 g-Cl/l the active chlorine
production rate levels off. Rudolf et al. (1995) reported that higher NaCl
concentrations decreased chlorate formation in divided electrolytic cell. The same was
observed by Czarnetzki and Janssen (1992) who performed their experiments in an
electrolytic cell divided by anionic membrane.
The effect of chloride concentration was tested with a Ti cathode and Ti/RuO2 anode,
both with dimensions of 9.15×5.1 cm2. The range of chloride concentrations in these
experiments was between 0.28 and 2.26 kA/m2 (10-80 g Cl/l). All experiments were
conducted with current density of 0.64 kA/m2, which was shown to result in the
highest current efficiency in the presence of 0.85 M Cl- (Subsection 4.2.3.3). The
oxidation rates were determined by linear regression of the –d[Fe(II)]/dt curve in the
interval of 2 to1 g Fe(II)/l. The results are presented in Fig. 34 and 35.
The voltage measured in the experiments decreased with the Cl- concentration, which
means that for higher Cl- concentration lower voltage is required to obtain the same
current density.
The increase in chloride ions concentration from 0.28 to 0.56 M increased
significantly both the oxidation rate and the current efficiency. At Cl- concentrations
higher than 0.85 M the effect was minor, similar to the results reported by Khelifa et
al. (2004). It is thus plausible to assume that at high Cl- concentrations the Fe(II)
concentration becomes the limiting factor of the Fe(II) oxidation reaction.
89
Fe(II) concentration [g Fe/l]
Figure 34. Concentration of Fe(II) as a function of time in the indirect electrochemical
4.5
72
4.0
64
3.5
56
3.0
48
2.5
40
2.0
32
1.5
24
-d[Fe(II)]/dt
voltage
1.0
16
Current efficiency
0.5
Current efficiency [%]
Voltage [V]
-d[Fe(II)]/dt [g Fe/h]
experiments with different chloride concentrations. Idensity=0.64 kA/m2
8
0.0
0
0.0
0.5
1.0
1.5
2.0
Concentration of Cl- [M]
2.5
Figure 35. Effect of Cl- concentration on Fe(II) oxidation rate, voltage and current
efficiency
90
4.3.3.5. Potential chlorine loss
Chlorine escape during the electrolysis was measured in the experiments that were
conducted with varied current densities and Cl- concentraion of 0.85 M (Subsection
4.2.3.3). Chlorine escape was observed only in the experiments with current densities
of 0.64 and 1.71 kA/m2. The amounts of Cl2(g) lost were 0.2 and 0.4 mM respectively,
less than 0.06% of total Cl- concentration. It should also be noted that both
experiments lasted much longer the point where all the Fe(II) was oxidized. Industrial
scale reactors should however take this phenomenon into consideration.
4.3.3.6. Indirect electrooxidation - energy cost
The highest current efficiency (59%) was found at current density of 0.64 kA/m2 in
the presence of 30 g Cl-/l (0.85 M). The voltage was 2.87 V and the oxidation rate was
3.7 g Fe(II)/h, which is equivalent to 1.13 g H2S/h (three orders of magnitude higher
than the rate obtained in the catalytic oxidation). The energy needed under such
conditions is 7.6×10-3 kW·h normalized for the removal of 1 gram of H2S(g),
calculated according to Eq. (61).
(
2.87 V × 0.64 kA
(
))
× 9.15 × 5.1 cm 2
kW ⋅ h
m2
= 7.6 * 10 −3
gH S
g H2S
1.13 2
h
(61)
Assuming that the cost of electricity is 0.12 $/kW·h the cost of removing 1 kg of
H2S(g) would be 0.91 $, which does not appear excessive.
7.6 * 10 − 3
kW ⋅ h
$
$
$
× 0.12
= 9.1 * 10 − 4
= 0.91
g H 2S
kW ⋅ h
g H 2S
kg H 2 S
(62)
1kg of H2S corresponds, for example, to 1455 m3 of air contaminated with 500 ppm of
H2S (a relatively high concentration).
mg H 2 S
500 ppm × 1.375
m 3 air × 1455m 3 = 1kg
air
H 2S
ppm H 2 S
(63)
4.3.3.7. Volume of solution in the electrolytic reactor
An electrolytic cell that is designed to oxidize, for example, 3.7 g Fe(II)/h, can
remove 1.13 g H2S/h. The amount of 1.13 g H2S corresponds to 1636 l of air
contaminated with 500 ppm of H2S, which can be treated by a recycle flow rate of
27.3 l/min. If the flux of air is set at 1.1 m3air/min/m3solution the volume of the
91
absorbing solution should be about 25 l. The volume of the electrolytic cell used in
the experiments reported in the current work was about 60 ml.
1.13
g H 2S
h
× 1455
l air
g H 2S
1.1l air / min/ l solution
1636
=
l air
h
l air
min
= 24.8l solution (64)
/ min/ l solution
27.3
1.1l air / min/ l solution × 60 min
=
h
1.1l air
4.3.3.8. Electrode material
Titanium, which was used as the cathode in the electrochemical oxidation
experiments, is known to have a high corrosion resistance in active chlorine solutions.
At acidic conditions (pH<5.0), when cathodic polarization is applied, the formation of
titanium hydride (TiH2) was reported to occur (Videm et al., 2008), which leads to the
embrittlement of titanium. Consequently the use of titanium as hydrogen evolution
cathode under acidic conditions is restricted in industrial long term operations. An
alternative cathode material was not attempted.
4.3.3.9. Changes in pH
An increase in pH from pH1.0 to pH1.2 was observed in most experiments, due to H2
evolution on the cathode. Rudolf et al. (1995) observed that the pH of an electrolyzed
NaCl solution increased during electrolysis from an initial pH6 to pH9.0-9.2 after
several minutes finally stabilizing at the range pH9.55±0.15 after another 20 minutes.
From the standpoint of the total process this increase in pH should not be a problem,
since the oxidation of H2S to S0 increases H+ concentration in the solution at an equal
rate.
To summarize all the above findings, indirect electro-oxidation of Fe(II) was found
much more suitable than direct oxidation for the H2S(g) removal process at pH1.0, due
to higher Fe(II) oxidation rates and significantly higher efficiency of H2S(g) reactiveabsorption in the presence of a high chloride concentration. The use of 30 g Cl-/l (0.85
M) seems to be the most advantageous chloride ion concentration, both from the
electrochemical oxidation perspective and from the H2S(g) reactive-absorption
standpoint.
92
10. Conclusions
•
Catalytic oxidation of Fe(II) by oxygen in the presence of copper and phosphate is
possible but not applicable for an LRSR process operated at pH1.0, since the
Cu(II) will precipitate with H2S and phosphate will precipitate with Fe(III).
•
The accelerating effect of phosphate on Fe(II) oxidation rate at pH1.0 is probably
related to the concentration of the H2PO4- ion.
•
The accelerating effect of copper on the Fe(II) oxidation rate at pH1.0 increases in
the presence of a higher O2 concentration and decreases in the presence of a
higher Fe(III) concentration.
•
The presence of Fe(III) at high concentrations decreases the Fe(II) oxidation rate,
both by O2 and electrochemically.
•
The reactive absorption efficiency of H2S(g) increases in the presence of chlorides,
although the addition of more than 30 g Cl/l seems to be superfluous. It is possible
that the positive effect of Cl- is due to high oxidation rate of H2S(aq) by the FeCl2+
complex.
•
Direct electrochemical oxidation of Fe(II) is not a viable option for an LRSR
process at pH1.0, mainly because the reactive-absorption efficiency of H2S(g) is
very low in the absence of chlorides.
•
Increase in current density (0.43 to 1.71 kA/m2) increases the rate of indirect
electro-oxidation of Fe(II), although the highest current efficiency (59%) was
found with current density of 0.64 kA/m2.
•
Increase in Cl- concentration up to 0.85 M (30 g Cl/l) increases the rate of indirect
electro-oxidation of Fe(II). Above 0.85 M the oxidation rate becomes
approximately constant.
•
Indirect electrochemical oxidation of Fe(II) is applicable for an LRSR process at
pH1.0: the Fe(II) oxidation rates are very high, the H2S(g) reactive-absorption
efficiency is high and precipitation is minimal even after 5 months.
93
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