APPLICATIONS OF FENTON AND FENTON-LIKE REACTIONS WITH
SUBSEQUENT HYDROXIDE PRECIPITATION FOR DERUSTING
WASTEWATER TREATMENT
PISETH SOM
A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE
REQUIREMENTS FOR THE MASTER DEGREE OF ENGINEERING
IN CHEMICAL AND ENVIRONMENTAL ENGINEERING
THE FACULTY OF ENGINEERING
BURAPHA UNIVERSITY
MAY 2014
COPYRIGHT OF BURAPHA UNIVERSITY
ii
CONTENTS
Page
ABSTRACT ...............................................................................................................
CONTENTS ...................................................................................................................ii
LIST OF TABLES ......................................................................................................... v
LIST OF FIGURES ...................................................................................................... vi
ABBREVIATION.......................................................................................................viii
CHAPTER 1 INTRODUCTION ................................................................................... 1
Statements and Significant of Problems .................................................................. 1
Objectives ................................................................................................................ 3
Research Hypothesis ................................................................................................ 4
Scope of the Study ................................................................................................... 4
Significance of the Study ......................................................................................... 4
CHAPTER 2 LITERATURE REVIEW ....................................................................... 6
Advanced Oxidation Processes (AOPs)................................................................... 6
Fenton’s Reagent and Reaction Mechanism ............................................................ 8
Basic Principle ................................................................................................ 8
Fenton Reaction .............................................................................................. 9
Hydroxyl Radical Reaction with Organic Compounds ......................................... 11
Iron Ligand, Chelators and Coordination .............................................................. 13
Factors Affecting Fenton and Fenton-like Process ................................................ 14
Effect of pH................................................................................................... 14
Effect of Temperature ................................................................................... 15
Effect of Iron Concentration ......................................................................... 15
Effect of H2O2 Concentration ........................................................................ 16
Effect of Reaction Time ................................................................................ 18
Chelating Agents Degradation by Various Fenton Processes ................................ 19
CHAPTER 3 RESEARCH METHODOLOGY .......................................................... 25
Derusting Wastewater Characteristics ................................................................... 25
Materials and Chemical Reagents .......................................................................... 25
Experimental Design and Procedure ...................................................................... 27
iii
Determine wastewater characteristics ........................................................... 27
Hydroxide Precipitation of Iron Before Fenton and Fenton-like Processes . 28
Effects of Initial pH on Fenton-like Process ................................................. 29
Effects of H2O2 Concentration on Fenton-like Process................................. 30
Effects of Reaction Time on Fenton-like Process ......................................... 31
Hydroxide Precipitation of Iron After Fenton-like Process .......................... 32
Effects of Initial pH on Fenton Process ........................................................ 32
Effects of Fe2+ Concentration on Fenton Process ......................................... 33
Effects of H2O2 Concentration on Fenton Process........................................ 34
Effects of Reaction Time on Fenton Process ................................................ 35
Hydroxide Precipitation of Iron After Fenton Process ................................. 36
Optimum Conditions .............................................................................................. 38
Analytical Method ................................................................................................. 39
Kinetic Study ......................................................................................................... 39
CHAPTER 4 RESULTS AND DISCUSSION ............................................................ 24
Wastewater Characterization ................................................................................. 24
Hydroxide precipitation of iron before Fenton and Fenton-like reactions ............ 26
Fenton-like reaction ............................................................................................... 28
Effect of initial pH ........................................................................................ 29
Effect of H2O2 concentration ........................................................................ 33
Effect of reaction time................................................................................... 37
Effect of precipitation pH ............................................................................. 40
TDS and Conductivity Content after Fenton-like reaction ........................... 44
Ammonia nitrogen, Nitrite, Nitrate removal ................................................. 45
Fenton reaction with subsequent hydroxide precipitation ..................................... 52
Effect of Initial pH ........................................................................................ 52
Effect of Fe2+ concentration .......................................................................... 56
Effect of H2O2: Fe2+ molar ratio ................................................................... 61
Effect of Reaction Time ................................................................................ 65
Effect of precipitation pH ............................................................................. 68
TDS and Conductivity Content after Fenton reaction .................................. 74
Ammonium nitrogen, Nitrite, Nitrate removal ............................................. 75
Comparison between Fenton and Fenton-like reactions ........................................ 82
Kinetics of Fenton and Fenton-like reactions ........ Error! Bookmark not defined.
Quality of treated water and possible use .............. Error! Bookmark not defined.
iv
CHAPTER 5 CONCLUSION...................................................................................... 83
REFERENCES ............................................................................................................ 84
APPENDICES ............................................................................................................. 92
Appendix A: Experimental data............................................................................. 92
Appendix B: Chemical analysis procedures .......................................................... 92
Appendix C: Fenton’s reagent preparation ............................................................ 97
Appendix D: Conceptual experimental pictures .................................................... 99
v
LIST OF TABLES
Table
Page
Table 2.1 Oxidizing potential for conventional oxidizing agents .................................. 6
Table 2.2 Summary of Fenton process for various EDTA complex wastewaters ....... 22
vi
LIST OF FIGURES
Figure
Page
Figure 2.1 Classification of advanced oxidation processes (AOPs). ............................ 8
Figure 3.1 Hydroxide precipitation before Fenton and Fenton-like reactions ............ 29
Figure 3.2 Experimental procedure for Fenton reaction (Adding Fe2+) and
Fenton-like reaction (Without adding Fe2+) ............................................... 38
Figure 4-1 Effect of initial pH on removal efficiencies of (a) TCOD and (b) SCOD at
[H2O2] of 2M, reaction time of 60 min, and precipitation pH 8
31
Figure 4-2 Effect of initial pH on removal efficiency of (a) total iron and (b)
dissolved iron at [H2O2] of 2M, reaction time of 60 min, precipitation pH 8 ............. 32
Figure 4-3 Effect of reaction time (min) on removal efficiencies of (a) TCOD and (b)
SCOD at initial pH of 3, [H2O2] of 2.5 M and precipitation pH 8............................... 38
Figure 4-4 Removal efficiencies (R %) of (a) total iron and (b) dissolved iron at
initial pH of 3, [H2O2] of 2.5 M, and precipitation pH 8 ............................................. 39
Figure 4-5 Removal efficiencies (R %) of TCOD and SCOD by (a) Fenton-like
reaction (b) precipitation at initial pH of 3, [H2O2] of 2.5 M, and precipitation pH 8 41
Figure 4-6 Removal efficiencies (R %) of total iron and dissolved iron by (a) Fentonlike reaction (b) precipitation at initial pH of 3, [H2O2] of 2.5 M, and precipitation pH
8.................................................................................................................................... 42
Figure 4-7 Effects of initial pH on TDS and conductivity at [H2O2] of 2M, reaction
time of 60 min, and precipitation pH 8 ........................................................................ 44
Figure 4-8 Removal efficiencies of ammonium and nitrate by (a) Fenton-like reaction
(b) precipitation at optimum condition ........................................................................ 48
Figure 4-9 Effects of initial pH on removal efficiencies (R %) of TCOD and SCOD
by (a) Fenton reaction (b) precipitation at [Fe2+] of 0.05M, [H2O2] of 2 M, and
precipitation pH 8 ........................................................................................................ 54
Figure 4-10 Effects of initial pH on removal efficiencies (R %) of total iron and
dissolved iron by (a) Fenton reaction (b) precipitation at [Fe2+] of 0.05M, [H2O2] of 2
M, and precipitation pH 8 ............................................................................................ 55
Figure 4-11 typical pH profile of Fenton reaction ...... Error! Bookmark not defined.
vii
Figure 4-12 Effects of initial Fe2+ concentration on removal efficiencies (R %) of
TCOD and SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [H2O2] of
2 M, and precipitation pH 8 ......................................................................................... 59
Figure 4-13 Effects of initial Fe2+ concentration on removal efficiencies (R %) of
total iron and dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of 3,
[H2O2] of 2 M, and precipitation pH 8 ......................................................................... 60
Figure 4-14 Effects of H2O2:Fe2+ molar ratio on removal efficiencies (R %) of TCOD
and SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of 0.05M,
and precipitation pH 8 .................................................................................................. 63
Figure 4-15 Effects of H2O2:Fe2+ molar ratio on removal efficiencies (R %) of total
iron and dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of 3,
[Fe2+] of 0.05 M, and precipitation pH 8 ..................................................................... 64
Figure 4-16 Effects of reaction time on removal efficiencies (R %) of TCOD and
SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of 0.05 M,
H2O2:Fe2+ molar ratio of 40, and precipitation pH 8.................................................... 66
Figure 4-17 Effects of reaction time on removal efficiencies (R %) of total iron and
dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of 0.05
M, H2O2:Fe2+ molar ratio of 40, and precipitation pH 8 .............................................. 67
Figure 4-18 Effects of precipitation pH on removal efficiencies (R %) of TCOD and
SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of 0.05 M,
H2O2:Fe2+ molar ratio of 40 and reaction time of 20 min ............................................ 71
Figure 4-19 Effects of precipitation pH on removal efficiencies (R %) of total iron
and dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of
0.05 M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20 min............................... 72
Figure 4-20 Concentration of total iron, Fe2+, and Fe3+ in (a) Fenton reaction effluent
(b) precipitation effluent at different precipitation pH at initial pH of 3, [Fe2+] of 0.05
M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20 min....................................... 73
Figure 4-21 Effects of initial pH on TDS and conductivity contents at [H2O2] of 2M,
reaction time of 60 min, and precipitation pH 8 .......................................................... 74
Figure 4-22 Removal efficiencies of ammonium and nitrate by (a) Fenton reaction (b)
precipitation at optimum condition pH at initial pH of 3, [Fe2+] of 0.05 M, H2O2:Fe2+
molar ratio of 40 and reaction time of 20 min ............................................................. 78
viii
ABBREVIATION
AOPs
: Advanced Oxidation Processes
EDTA
: ethylenediamine tetraacetic acid
COD
: chemical oxygen demand
BOD5
: biological oxygen demand in 5 day
[H3O2]+ : oxonium ions
H2O2
: hydrogen peroxide
HO2•
: perhydroxyl radical
OH•
: hydroxyl radical
Fe3O4
: magnetite
Fe2+
: Ferrous ion
Fe3+
: Ferric ion
EOP
: electrochemical oxidation potential
1
CHAPTER 1
INTRODUCTION
This chapter covers the fundamental background of research and problems
in consideration of Advanced Oxidation Processes (AOPs) based on Fenton and
Fenton-like processes for derusting wastewater; then, the research objectives and
research hypothesis are formulated accordingly. Finally, significance and scope of the
study are also provided.
Statements and Significant of Problems
Chemical cleaning of pipes, tanks, boilers, and power plants has been
operated to remove the deposits and scales for reactivation and reuse of them. There
are various types of chemicals that have been used for cleaning depending on the
equipment including inorganic acids, organic acids, chelating agents, alkali agents and
aids agents (Pliego et al., 2013). The inorganic acids include hydrochloric acid,
sulfuric acid and nitric acid. The hydrochloric acid is the most widely used for
chemical cleaning. The examples of organic acids are citric acid, glycolic acid, and
formic acid. The organic acids are used extensively for cleaning of recent new boilers.
The most widely used chelating agent is ethylenediamine tetraacetic acid (EDTA).
The ammonia, which is alkali agent, is used to clean the scale containing large
quantities of copper. The aids agents such as acid inhibitors and reducing agents are
used to reduce and to prevent the corrosion of the materials, respectively. The sodium
nitrite can be used as inhibitor for protection of carbon steel in salt solution (Hayyan,
et al., 2012). During the cleaning operation, two methods for dissolving encrustation
or rust are applied. First method is a two-step process: first stage uses inhibited
hydrochloric acid solution for iron oxide dissolution followed by the second stage of
dissolving the metallic copper by ammonia and oxidizing agents. Another method
involves a single cleaning stage. In this method, iron oxide and metallic copper are
dissolved simultaneously by using hydrochloric acid in the presence of chelating
agents and citric acid. Consequently, the cleaning wastewater often contains large
amounts of iron and copper including high concentration of chelating agents (Huang
2
et al., 2000; Bansal, 2012). Iron (Fe3+) is the most prevalent cation, generally present
at a concentration of 1000-10000 mg/L. Copper is the second most abundant metal
with minor level of nickel, chromium, and zinc, typically, present at the concentration
less than 100 mg/L (Huang et al., 2000; Kim et al., 2010).
The EDTA and citric acid are used at the concentrations of 2-5% and up to
10% by weight, respectively, in cleaning process (Huang et al., 2000; Kim et al.,
2010). Chelating heavy mental wastewater must be treated not only for the toxic
heavy mental, but also the chelating agents. Heavy metals are considered toxic to
human being and aquatic life. Furthermore, the EDTA causes the complexation and
mobilization of heavy metals. The EDTA complexation is biologically persistent and
cannot be readily degraded by conventional biological treatment processes (Ghiselli et
al., 2004; Citra et al., 2011). The presence of chelated complex causes constraints and
ineffective application of lime or caustic treatment, chemical precipitation, ion
exchange as reported in the literatures (Citra et al., 2011; Lan et al., 2012; Fu et al,
2009). Metal chelated wastewater can be treated by electrochemical reduction (Huang
et al., 2000) and interior microelectrolysis (Lan et al., 2012). Both processes can
successfully remove metal; however, interior microelectrolysis cannot remove or
degrade chelating EDTA and electrochemical reduction can achieve EDTA recovery
for reuse. To remove metal and mineralize the metal-EDTA complexes, there is an
urgent need to search for a feasible, efficient, economical, and eco-friendly
approach(Bautista et al., 2008; Bianco et al., 2011).
For last few decades, advanced oxidation processes (AOPs) are known for
their capability to mineralize, decompose, and degrade non-biodegradable organic
compounds (Poyatos et al., 2010; Ameta et al., 2012). Particularly, Fenton and
Fenton-like processes are adopted for wastewaters treatment in terms of organic
pollutant destruction, toxicity reduction, biodegradability improvement, COD
removal, odor and color removal, and heavy metal removal due to the economic
advantages, ease of application, and effectiveness (Matthew Tarr, 2003; Bautista et
al., 2008a; Lucas & Peres, 2009; Bianco et al., 2011). Fenton reaction is one of the
AOPs that has been commonly applied for industrial wastewater including textile
effluent (Kang et al., 2002; Karthikeyan et al., 2011), olive oil effluent (Lucas &
Peres, 2009; Kiril Mert et al., 2010), pulp and paper mill effluent (Pirkanniemi et al.,
3
2007), cosmetic wastewater (Bautista et al., 2007), bleaching effluent (Wang et al.,
2011), highly polluted industrial wastewater (San Sebastián Martinez et al., 2003),
complex industrial wastewater (Bianco et al., 2011). However, application of Fenton
and Fenton-like reactions for boilers chemical cleaning wastewater is not extensively
documented. An integration of Fenton oxidation with other conventional treatment
methods have been conducted to degrade EDTA complex and to removal metals from
the waste stream. Synthetic NiEDTA was successfully removed using Fenton and
Fenton-like reactions followed by precipitation (Fu et al., 2009, 2012). The
degradation of Cu-EDTA complex can be achieved (Lan, et al., 2012) with interior
microelectrolysis and Fenton oxidation–coagulation. However, Fe-EDTA complex
has not been conducted yet. To our knowledge; therefore, Fe-EDTA removal by
Fenton and Fenton-like is important since their applications are limited.
As mentioned above, chemical cleaning wastewater contains high iron
species including Fe2O3,/Fe3O4 (rust) and Fe2+ or Fe3+ depending on the pH, which
can reach up to hundreds of mg/L. It is assumed that Fenton or Fenton-like reactions
should take place to generate hydroxyl radicals (OH•) when H2O2 is added to the ironrich wastewater because the rust (Fe2O3/Fe3O4) particles and iron (Fe2+/Fe3+), which
have already presented in wastewater, could be effective catalysts in the generation of
strong oxidant (Kitis & Kaplan, 2007; Kim et al., 2010; Lan et al., 2012).
Objectives
The overall objective of this study was to evaluate the feasibility and
efficiency of Fenton and Fenton-like oxidations for removal of organic pollutants
measured as COD and inorganic pollutants including various iron species
concentrations as the main parameters and chemicals in derusting wastewater.
Following specific objectives were included:
1. To determine the optimum initial parameters of Fenton and Fenton-like
reactions including pH, Fe2+ concentration, H2O2 concentration for the
treatment of derusting industrial wastewater.
2. To determine the optimum reaction time and reaction kinetics for the
treatment of derusting industrial wastewater
4
3. To determine the optimum precipitation pH for Fenton and Fenton-like
reactions for the treatment of derusting industrial wastewater.
4. To investigate the effects of Fenton and Fenton-like reactions on the
ammonia and nitrite removals in the treatment of derusting industrial
wastewater.
Research Hypothesis
1. The presence of chelating agent, EDTA, can inhibit the precipitation of iron
in the derusting wastewater.
2. Utilization of existing Fe (III)/Fe (II) and additional iron can be beneficial for
Fenton and Fenton-like oxidation for degradation of EDTA complex in term
of COD reduction and total Fe removal.
Scope of the Study
This study was limited with following conditions.
1. Treatment performance evaluation was conducted using Jar test apparatus
under normal laboratory room temperature at Department of Chemical
Engineering, Faculty of Engineering, Burapha University.
2. Real wastewater taken from Kation Power Ltd (Thailand) was used
throughout the experiments
3. Organic degradation was measured in chemical oxygen demand (COD)
4. Oxidation products were not investigated in this study
Significance of the Study
The results of this study can provide the following contributions:
Firstly, this study demonstrates the fessibility of Fenton and Fenton-like
process applications as methods to solve the encountered derusting wastewater
treatment problem as practiced in accordance with standard effluent stipulated in
national regulation.
Secondly, even though Fenton oxidation have been applied extensively and
enormously in many differrent types of wastewater, its application for derusting
wastewater was not well documented in literatures. Thus, this study will contribute to
5
comprehensive and extensive knowlegde and discussion on real wastewater treatment
which is known to be contiminated with chelating organic compounds and high metal
concentration.
Finally, it is probably advantageous for Fenton and Fenton-like process to
utilize iron metals (fersric and ferrous ions) that have already existed in cleaning
wastewater. If they do, there will be economical and cost-effective for reagents usages
for the treatment of this wastewater.
6
CHAPTER 2
LITERATURE REVIEW
This chapter provides a comprehensive review on advanced oxidation
processes (AOPs). Next, theoretical and empirical reviews on Fenton and Fenton-like
reaction mechanisms influencing factors and their applications were conducted.
Finally, the applications of Fenton oxidation for chelating agent, EDTA, were also
reviewed.
Advanced Oxidation Processes (AOPs)
The development of cost-effective technical solutions isneeded to deal with
the increasingly complex problems arising in the field of industrial wastewater.
Recently, advanced oxidation processes (AOPs) have been applied successfully for
the removal or degradation of recalcitrant pollutants based on the high oxidative
power of the hydroxyl radical (HO•). It has electrochemical oxidation potential (EOP)
of 2.8 V, which is comparatively be second to fluorine as shown in table 2.1 (Poyatos
et al., 2010).
Table 2-1 Oxidizing potential for conventional oxidizing agents
Oxidizing agent
Oxidation Potential (EOP), V
EOP relative to
Chlorine (V)
Fluorine
3.06
2.25
Hydroxyl radical (HO•)
2.80
2.05
Oxygen (atomic)
2.42
1.78
Ozone
2.08
1.52
Hydrogen peroxide
1.78
1.30
Hypochlorite
1.49
1.10
Chlorine
1.36
1.00
Chlorine dioxide
1.27
0.93
Oxygen (molecular)
1.23
0.90
7
Source: Poyatos et al., (2010)
A chemical wastewater treatment using AOPs can produce the complete
mineralization of pollutants to CO2, water, and inorganic compounds, or at least their
transformation into more harmless products. Furthermore, the partial decomposition
of non-biodegradable organic pollutants can lead to biodegradable intermediates;
therefore, AOPs are commonly applied as pre-treatments processes, followed by
biological or chemical processes (Poyatos et al., 2010). AOPs represent the newest
methods in H2O2 technology which include photochemical degradation processes
(UV/O3, UV/ H2O2), photocatalysis (TiO2/UV, photo-Fenton reaction), and chemical
oxidation processes (O3, O3/H2O2, H2O2/Fe2+). Although advanced oxidation
processes (AOPs) have employed different reagent systems, they all produce hydroxyl
radicals. These radicals are very reactive and they can attack most organic
compounds nonselectively (Kalra et al., 2011; Lucas & Peres, 2009; Poyatos et al.,
2010). Advanced oxidation processes (AOPs) can be classified either as homogeneous
or heterogeneous. Homogeneous processes can be further subdivided into energyactivated and non-energy activated processes as shown in Figure. 2.1. The following
sections describe a wide range of advanced oxidation systems that are currently being
studied for their possible use in wastewater treatment (Poyatos et al., 2010).
Among advanced oxidation technologies, Fenton oxidation has been
frequently involved in many different industrial wastewater treatment processes for
degrading and remediating of a wide range of contaminants, predominately toxic,
recalcitrant, and persistent organic pollutants (POPs). It is also due to economic
advantages, ease of application, and effectiveness in the contaminant reduction and
mineralization (Matthew Tarr, 2003). It was also considered that Fenton oxidation
presents one of the best methods for clean and safe processes for the degradation of
organics even at higher initial organic content (Bianco et al., 2011; Lucas & Peres,
2009).
8
Advanced Oxidation Processes
Homogeneuos
process
Using Energy
Ultraviolet
Radiation
- O3/UV
- H2O2/UV
- H2O2/O3/UV
- PhotoFenton(Fe2+/
H2O2/UV)
Ultrasound
Energy
- O3/US
- O3/US
Heterogeneuos
process
Without Energy
Electrical
Energy
- Anodic
Oxidation
- Electro-Fenton
- O3 in alkaline
Medium
- O3/ H2O2
- Fenton Process
Fe2+/ H2O2
- Fentton-like
Fe3+/H2O2
Fe0/H2O2
- Catalytic
Ozonization
- Photocatalytic
Ozonization
-Heterogeneous
Photo-catalysis
Figure 2-1 Classification of advanced oxidation processes (AOPs).
Fenton’s Reagent and Reaction Mechanism
Basic Principle
The term Fenton’s reagent refers to the aqueous mixture of Fe (II) and
hydrogen peroxide. The Fenton’s reagent was first discovered and used by H. J. H.
Fenton in 1894 when he observed that the rate of oxidation of tartaric acid increased
dramatically when dilute hydrogen peroxide with the solution containing dissolved
Fe2+ions. Forty years later, after a controversial history about the reaction mechanism
of Fenton’s reaction, its reaction mechanism was interpreted by Haber and Weiss in
1934 that Fenton’s chemistry is a reaction between hydrogen peroxide (H2O2) and
Fe2+ ions forming hydroxyl radicals, which is the main oxidizing agent. However the
hydroxyl radical mechanism of the Fenton’s reaction for toxic organics degradation
was not applied until the late 1960s (Ciambelli et al., 2008; Matthew Tarr, 2003;
Neyens & Baeyens, 2003).
9
Fenton Reaction
The oxidation mechanism in the Fenton process involves ferrous ions (Fe2+)
to react with hydrogen peroxide, producing hydroxyl radicals with powerful oxidizing
ability to degrade organic pollutants. The oxidation mechanism of Fenton reaction is
very complex, but the widely accepted major chemical reactions are summarized as
shown below (Ameta et al., 2012; Bianco et al., 2011; Jiang et al., 2010; Lee &
Shoda, 2008; Lucas & Peres, 2009; Matthew Tarr, 2003; Neyens & Baeyens, 2003;
Munter, 2001).
Fe2+ + H2O2→ Fe3+ + OH• + OH−
k = 70 M-1s-1
(2-1)
k =107 -1010 M-1s-1
(2-2)
R• + Fe3+ → R+ + Fe2+
-
(2-3)
Fe2++ OH• → Fe3++ OH−
k = 3.2 108 M-1s-1
(2-4)
H2O2 + OH•→ HO2• + H2O
k =3.3 107 M-1s-1
(2-5)
RH + OH• → R•+ H2O
As shown in equation (2.1), the ferrous iron (Fe2+ ) initiates and catalyses the
decomposition of hydrogen peroxide (H2O2) to generate the hydroxyl radicals (OH•).
The reaction (2.1) is commonly known as the main reaction of Fenton process
(Neyens & Baeyens, 2003). The generated hydroxyl radical reacts immediately with
organic substances (RH) resulting in a free organic radicals (R•). These radicals are
subsequently oxidized by ferric ion to generate other oxidation products (Matthew
Tarr, 2003). In addition to the main reaction, various additional competitive or
scavenging reactions are also possible involving ferrous ions (Fe2+), hydroxyl radicals
(OH•), and hydrogen peroxide (H2O2) as listed in reactions (2.4)-(2.5). During the
reaction, the newly formed ferric ions (Fe3+) may continuously catalyze hydrogen
peroxide to produce ferrous ions and perhydroxyl radical (HO2•). The reaction of
hydrogen peroxide with ferric ions is referred to Fenton-like reaction (Ameta et al.,
2012; Bianco et al., 2011; Matthew Tarr, 2003; Neyens & Baeyens, 2003a). Fentonlike reactions are listed as below:
Fe3++ H2O2→ Fe2++ H++ HO2•
k = 0.001-0.01 M-1s-1
(2-6)
10
Fe3+ + HO2•→ Fe2++ H+ + O2
Fe2+ + H2O2→ Fe3+ + OH• + OH−
RH +OH•→ R•+ H2O
k = 1.2 106 M-1s-1
(2-7)
k = 70 M-1s-1
(2-8)
K = 107 -1010 M-1s-1
(2-9)
In the presence of organic substrates (RH), highly reactive hydroxyl radical
which is species with a relatively short life-span (rate constants in the range 107 -1010
M-1s-1), undergoes oxidation generating a new radical (R•) as shown in reaction (2.9).
The possible organic compounds present in reaction mixture can suffer an abstraction
of a hydrogen atom (proton abstraction) or addition of hydroxyl radical (OH•) with the
production of organic radicals (R•) which can subsequently be oxidized by ferric ions
(Fe3+) as indicated in reaction (2.3). Indeed, the reaction (2.3) regenerates ferrous ions
(Fe2+) which ensure the continuity of the chain reaction. As long as the concentration
of reactants are not limited or available in the system, the iron species continually
cycle between Fe2+ and Fe3+ unless additional reaction result in formation of insoluble
iron oxides and hydroxides. This can lead ultimately to the decomposition of organic
substrate in carbon dioxide (CO2) and water inorganic salts (Lucas & Peres, 2009;
Matthew Tarr, 2003; Neyens & Baeyens, 2003).
The conventional Fenton has been modified to improve treatment efficiency
with the reduced inorganic sludge production and prevention of inhibition reaction of
some ions. Those modified Fenton technologies includes photo-Fenton, electroFenton, electro-photo Fenton and Fenton-like reaction. Fenton-like process uses other
transition metal catalyst other than Fe2+ (Fu et al., 2009). The conventional Fenton has
been applied numerously while Fenton-like is not well elucidated. The introduction of
lower cost Fe3+ in Fenton-like process may overcome the drawback of conventional
Fenton (S. Wang, 2008).
Recent applications of other transition metals in addition to Fe2+ including
Fe-containing zeolites, soluble manganese (II) and amorphous and crystalline
manganes (IV) oxide, soluble Fe3+, mixture of Fe2+/Cu2+ and Fe3+/Cu2+, suspended
iron powder, clay-based Fe nanocomposite and zero valent iron (ZVI) were
investigated. However, ZVI and Fe3+ have been commonly used as catalysts in
Fenton-like reaction due to their comparable efficiency and capacity (Fu et al., 2009&
2013; Hodaifa et al., 2013; Jiang et al., 2010 & 2013). Since the Fenton-like reaction
11
can be applied interchangeably and comparatively with Fenton reaction, it was
recently selected for wastewater treatment application in term of cost-effectiveness,
efficiency, and easy of application (Fu et al., 2009; Hodaifa et al., 2013; Jiang et al.,
2010, 2013; Kim et al., 2010; Kiril Mert et al., 2010; Li et al., 2013). Other
investigations of Fenton and Fenton-like process by using iron originated in
wastewater still remain questionable even though iron waste existed in the wastewater
was feasibly use as catalyst for Fenton reaction (Lan et al., (2012).
Jaing et al. (2013) has indicated the interconversion of Fe(III)/Fe(II) in
Fenton and Fenton-like reaction that they are co-occurring or coexisting. A Fentonlike reaction involves a classical Fenton reaction, and Fenton reaction may involve a
Fenton-like reaction step. However, Jaing et al. (2010) and Neyens & Baeyens (2003)
demonstrated conventional Fenton reaction was referred to the Fe2+/H2O2 system,
whereas Fenton-like reaction was included in the Fe3+/H2O2 system. Therefore, the
reaction mechanisms are similar in both systems, but are different in terms of catalysts
that are utilized to initiate the reaction.
Hydroxyl Radical Reaction with Organic Compounds
For the reaction of hydroxyl radical with organic species, there are three
common reaction pathways: (a) hydroxyl radical addition to an unsaturated compound
(aromatic or aliphatic) to form the free radical products, (b) hydrogen abstraction
where an organic free radical and water are formed (c) electron transfer, where ions of
higher valence state are formed reducing hydroxyl radical to hydroxide ions (Matthew
Tarr, 2003; Munter, 2001; Neyens & Baeyens, 2003a). Reaction pathways are shown
below:
RH + OH•→ (OH)RH•
(Hydroxyl Radical
C6H6 + OH•→ (OH)C6H6•
Addition)
RH + OH• → R• + H2O
CH3OH + OH• → CH2OH• + H2O
(Hydrogen Abstraction)
(2-10)
(2.11)
12
RH + OH• → (RH)• + + OH−
[Fe(CH)6]4− + OH• → [Fe(CH)6]3− + OH−
(Direct Electron Transfer)
(2-12)
Additional reactants including Fe2+, Fe3+, H2O, O2, H+ ,OH•, other metals,
other organics, and other radicals present in the system are necessary to complete
these subsequent reactions. Further oxidation processes continuously occur and
dimerizeation can also occur if the initially formed radical species reacts with another
identical radical. Other possible reactions including radical interaction where the
hydroxyl radical reacts with other hydroxyl radical to combine or to disproportionate
to form the stable products (Munter, 2001; Neyens & Baeyens, 2003).They are shown
as following:
OH• + OH• → H2O2 (dimerization of OH•)
(2-13)
R• + H2O2 → ROH + OH•
(2-14)
R• + O2 → ROO•
(2-15)
ROO• + RH → ROOH + R•
(2-16)
The organic free radical produced in the above reactions may then be
oxidized by Fe3+ reduced by Fe2+, or dimerized according to the following reactions.
R• + Fe3+ -oxidation → R+ + Fe2+
(2.17)
R• + Fe2+-reduction → R− + Fe3+
(2.18)
R• + R •-dimerization → R−R
(2.19)
By applying Fenton’s Reagent for industrial waste treatment, the
predominant reactions are hydrogen abstraction and oxygen addition. Typical rates of
reaction between the hydroxyl radical and organic materials are 109 – 1010 k (M-1 s-1)
(Matthew Tarr, 2003).
13
Iron Ligand, Chelators and Coordination
Chelating agents still remained contradicted for Fenton reaction. Addition or
presence of resolubilizing or chelating agents cause an increase in the occurrence of
reaction in the catalytic Fenton process. In contrast, chelating agents can interfere the
Fenton process by scavenging ability of the chelators. A good scavenger may appear
to have a lower production rate of hydroxyl radical due to rapid trapping of the radical
by the chelator. In addition, very strong iron chelators inhibit the formation of
hydroxyl radical. Iron ligands can also act as hydroxyl radical scavengers. Ligands are
more likely to react with hydroxyl radical than pollutants that are not in close
proximity to the iron because radical is always formed in close proximity to these
ligands. Such coordination will alter the kinetics of hydroxyl radical formation as well
as the dynamics of hydroxyl radical interaction with pollutants. Matthew Tarr(2003)
concluded that the inability of hydroxyl radical to reach sorbed or sequestered
pollutants is one of the major drawback to the application of Fenton degradation
method. However, it is suggested that aggressive conditions including high H2O2
concentration could make possibility for direct degradation of sorbed species.
Several studies have been investigated for the effect of chelators on Fenton
reaction. Addition of chelators to Fe(III)-H2O2 systems (Fenton-like reaction) allows
for effective degradation at near neutral pH values. The influence of the iron chelators
form increased solubility of iron species at higher pH value. Iron chelators improved
the Fenton oxidation of pollutant by increasing iron solubility and increases the rate
constant for hydroxyl radical formation from peroxide. The chelators also act as
hydroxyl radical scavengers from potential interaction with pollutants. The relative
efficiencies of the chelators for hydroxyl radical formation determine whether the
added chelators will have a positive or negative effect on radical formation. The
complexation of EDTA with iron minimized free ions for Fenton’s oxidation,
resulting in a slow generation of OH radical (Sillanpää et al., 2011). However, the
chelating agent may activate H2O2 oxidation at a neutral pH range. This pH ranges
might affect the Fenton’s process due to iron precipitation (Ghiselli et al., 2004). It
reaches to a conclusion that the presence of iron ligands and coordination could bring
both positive and negative influences on Fenton process depending on specific
property of iron-coordinating complex.
14
Factors Affecting Fenton and Fenton-like Process
The significant factors affecting both processes are H2O2 concentration and
iron concentrations, pH, reaction time, temperature and initial pollutant concentration.
Effect of pH
The suitable pH for Fenton process is also determined to be between 3 to 6
according to USperoxide (2012). However, different values of the operating pH have
also been reported (Matthew Tarr, 2003; Neyens & Baeyens, 2003a). According to
Wang et al. (2011), Fenton oxidation presented the maximum catalytic activity at pH
2.8-3.0. Similarly, Fu et al.(2009, 2012) and Lan et al.(2012) found the optimal pH of
3 and 2-5, accordingly for metal-EDTA complex wastewater treatment. A study on
EDTA degradation by Fenton process with pH ranged from 2 to 7 found that
degradation of EDTA decreased from 80.3% to 27.5% over the reaction time of 10
min and optimum pH range for Fenton oxidation was 2-4 (Lou & Huang, 2009). At
very low pH, H2O2 is stabilized as oxonium ions (H3O2+).The reaction between •OH
and H+ also occurs. Fe2+ regeneration by the reaction of Fe3+ with H2O2 is inhibited at
more acidic pH value (Wang et al., 2011). On the other hand, at high pH (pH > 3),
oxidation yield of the process decreases due to the precipitation of Fe3+ as Fe(OH)3
which hindered the reaction between Fe3+ and H2O2 and thus influenced the
regeneration of Fe2+. Moreover, Fe(OH)3 functionally catalyzes the decomposition of
H2O2 into O2 and H2O which decrease the production of hydroxyl radical (•OH) (Fu et
al., 2012, Wang et al,. 2012, Bautista et al., 2008).
A second aspect of pH deals with its shift as the reaction progresses. During
the Fenton reaction, an initial wastewater pH typically degreases. This pH decrease is
caused by the addition of FeSO4 catalyst which typically contains residual H2SO4.
more pronounced drop in pH occurs as the H2O2 is added, and continues gradually at
a rate which is largely dependent on catalyst concentration. This drop in pH is
attributed to the fragmenting of organic material into organic acids. Therefore, pH of
solution has to be controlled for Fenton reaction to ensure the reaction occurs
(USperoxide, 2012).
15
Effect of Temperature
The effect of temperature on the rate of reaction of the Fenton process
increases as the solution temperature increases. The application of temperature greater
than 40 °C, the treatment efficiency declined due to the decomposition of H2O2 into
oxygen and water. Fenton process has been normally conducted at temperature of 20
to 40°C(Bautista et al., 2008).A comparative study of Fenton and Fenton-like reaction
kinetics in decolorization of wastewater. The result has been indicated that
temperature had little influence on overall dye degradation in the range 15-45
°C(Wang, 2008). Dye degradation rate decreased when the temperature greater than
30 °C due to decomposition of H2O2 at higher temperature. Similarly, San Sebastián
Martinez et al., (2003) found that temperature showed only a mild positive effect on
COD removal. The significance of temperature influencing the Fenton and Fentonlike oxidation was clear that the increase of temperature could increase the removal
efficiency in the system because higher temperature increases the reaction between
hydrogen peroxide and Fe2+/Fe3+, and improve the generation rate of hydroxyl
radicals. The increase temperature from 25 to 50°C, the removal efficiency of Ni
increased from 72.1 to 97.2% for Fenton and from 74.3 to 96.7% for Fenton-like after
20 min (Fu et al. 2009). Since Fenton reaction is exothermic (optimal temperature
varied from 20 to 30 °C), it allow an industrial treatment of OMW without
temperature control (Nieto et al., 2011). Consequently, temperature was not
considered in the optimization of Fenton’s reaction in highly polluted industrial
wastewater. This leads to a conclusion that temperature is important but not necessary
for Fenton reactions because of exothermic effects of reaction leading to increase of
temperature in a suitable of range as found in the works of Bautista et al (2008);
Wang (2008); Fe et al. (2009, 2012);and Lan et al.(2012).
Effect of Iron Concentration
Iron concentration plays a vital role treatment efficiency of Fenton and
Fenton-like reactions because the production rate of hydroxyl radical (OH•) is
proportional to the concentration of iron and hydrogen peroxide. However, iron
content is the determining factors in sludge production as a challenge for Fenton
reaction (Wang et al., 2011). In the absence of iron, there is no evidence that OH• is
16
produced in wastewater. Inadequate concentration of iron in the operating condition
will lead to insufficient production of OH•, whereas overdosing of iron can favor the
scavenging reaction which prevents the reaction of OH• with contaminants resulting
in poor treatment efficiency (Matthew Tarr, 2003; Neyens & Baeyens, 2003). The
influence of ferrous concentration on EDTA degradation have been indicated that
increase of ferrous concentration from 10-4 M to 10-2 M resulting in the degradation of
EDTA from 29.8% to 98.5% at a reaction time of 10 min., respectively. However,
increasing Fe2+concentration from 10-2 M to 10-1 M decreased EDTA degradation
from 98.5% to 44.9%, accordingly. A higher Fe2+ dose provided the scavenging
reaction between Fe2+ and OH• (Lou & Huang, 2009). Another study found that the
increase of initial Fe2+ or Fe3+ from 0 to 1.0 mM resulting in the increasing of removal
efficiency remarkably. When Fe2+ or Fe3+ concentration was 1.0 mM, Fenton and
Fenton-like systems achieved 92.8% and 94.7% of Ni removal efficiencies after 60
min. of reaction time, accordingly. However, further increase of Fe2+ and Fe3+
concentration did not achieve the improvement in Ni removal (Fu et al., 2009). This
indicated that the use of much Fe2+concentration could lead to the self-scavenging of
OH• by Fe2+ as explained in literatures (Matthew Tarr, 2003; Neyens & Baeyens,
2003). A minimal threshold concentration of 3-15 mg/L Fe which allows the reaction
to proceed within a reasonable period regardless of the concentration of organic
materials. A constant ratio of Fe:substrate above the minimal threshold, typically 1
part Fe per 10-50 parts substrate, which produces the desired end products. The ratio
of Fe:substrate may affect the distribution of reaction products. A supplemental
aliquot of Fe which saturates the chelating properties in the wastewater; thereby,
availing unsequestered iron to catalyze the formation of hydroxyl radicals. Iron dose
may also be expressed as a ratio to H2O2 dose. Typical ranges are 1 part Fe2+ per 5-25
parts H2O2 (wt/wt) (USperoxide, 2012).
Effect of H2O2 Concentration
The amount of H2O2 is considered one of the most important factors in
Fenton and Fenton-like reaction owing to its economic cost, sources of OH•
generation, improvement of treatment efficiency and side effects in overdosing. The
17
H2O2 dose has to be fixed according to the initial pollutant concentration (Matthew
Tarr, 2003).
It is frequent to use an amount of H2O2 corresponding to the theoretical
stoichiometric H2O2 to chemical oxygen demand (COD) ratio, although it depends on
the response of the specific contaminants to oxidation and on the objective pursued in
term of reduction of the contaminant load (Neyens & Baeyens, 2003; Bautista et al.,
2007; Lan et al.,2012). Effect of H2O2 on the removal of COD was indicated that
increase in [H2O2]/[COD] from 0.5 to 2.0, the COD removal increased remarkably
from 73.6% to 89.4%. However, the further increase in [H2O2]/[COD] from 2.0 to 6.0,
the removal of COD was negligible or unchanged (Lan et al., 2012;Wang et al.,
2011). The marginal improvement of COD removal may be explained by the
scavenging effect of excessive H2O2 to OH• and recombination of OH• which were
supported in literatures (Neyens & Baeyens, 2003; Matthew Tarr, 2003; Bautista et
al., 2007, 2008; Wang, 2008; Wang et al., 2011; Lucas & Peres, 2009). Therefore,
stoichiometric relation between COD and H2O2 are significant for Fenton reaction and
acceptable [H2O2]/[COD] weight ratio should in the range of 2-4.
For most applications, it is important to optimize the molar ratio of
[Fe2+]/[H2O2] for estimation of reagent requirement and convenience of experiments
(Matthew Tarr, 2003; Neyens & Baeyens, 2003a). The presence of Fe2+ or Fe3+ salts
not only functions as catalytic reagents to decompose H2O2 for •OH generation, but
also reduces the scavenging effect of OH• radical from H2O2. The role of Fe3+ plays an
important role in oxidizing the target organic compound and producing OH• radical
through Fe2+ reaction (Kim et al., 2010). The [Fe2+/3+]/[H2O2] ratio is difficult to
specify and is varied according to the degradation of different pollutants covering the
range from 1:1 to 1:400 for a complete oxidation as reported in De Souza et al.(2006).
Effects of [Fe2+]/[H2O2] molar ratios of 1:50, 1:20, 1:10, 3:4 were conducted
for removal of initial COD of 300 mg/L by applying [H2O2]/[COD] of 4. Greater than
55% of COD removal was achieved in the first 10 min at higher [Fe2+]/[H2O2] molar
ratio. This results from higher generation of OH• radical according to reaction (2.1) as
shown previously. However, COD removal tended to decline in molar ratio of
[Fe2+]/[H2O2] greater than 1:20 due to quenching or scavenging effects of OH• radical
by excessive Fe2+ according to reaction (2.2). [Fe2+]/[H2O2] ratio of 1:20 attained
18
highest performance for greater than 85% of COD removal (Wang et al., 2011). To
achieve 90% removal of 362000 mg/L COD, it was required to maintain the optimal
[Fe2+]/[H2O2] molar ratio of 1:10, while [H2O2] was 3M (San Sebastián Martinez et
al., 2003). This molar ratio was comparatively found to be lower than that of
[Fe2+]/[H2O2] molar ratio at 1:15 resulting in the study of Lucas and Peres (2009). It is
clear that [Fe2+]/[H2O2] molar ratio varies according to type and concentration of
organic pollutant existing in the wastewater. The typical range of Fe2+]/[H2O2] ratios
are 1:5-25 as reported in Bautista et al. (2008) and USperoxide (2012).
Effect of Reaction Time
The time needed to complete a Fenton reaction depends on many variables
discussed above, most notably catalyst dose and wastewater strength. Typical reaction
times are 30-60 minutes for low strength wastewater. For more complex or more
concentrated wastes, the reaction may take several hours. Determination of reaction
completion prove troublesome (Matthew Tarr, 2003). A study on Fenton and Fentonlike reactions from 20–120 min. was conducted. Reaction time of 60 min for both
processes was determined for reduction of Ni concentration from 50 mg/L to 1 mg/L
and COD decreased from 252 mg/L to 53.3 mg/L, indicating about 78.8% COD
removal. After 60 min of reaction, the removal efficiency was marginal or almost
unchanged (Fu et al. 2009 & 2012). This reaction time for Fenton oxidation is
consistent with Lan et al. (2012), who found optimum reaction time at 60-80 min.
However, with heterogeneous and complicated characteristics of wastewater, it was
required 120 min for reduction of COD from 300 mg/L to 40 mg/L (Wang et al.,
2011).
The reaction time for a completion of Fenton reaction also depends on the its
reagents (Fe2+ and H2O2) because the contaminant degradation rate is proportional to
the hydroxyl radical produced (Matthew Tarr, 2003). San Sebastián Martinez et
al.(2003) and Jiang et al. (2013) achieved optimum efficiencies in the first 10 min of
Fenton reaction due to the fast reaction in the first stage of Fenton oxidation, while
prolonging the reaction time remained efficiency insignificantly changed. However, it
was required longer than an hour reaction time for metal-complex wastewater
treatment due the persistency of organic compounds (Pirkanniemi et al., 2003).
19
Therefore, the application of Fenton oxidation to industrial wastewater treatment
typically varies from 1 to 4 hours for optimal reaction time as reviewed in Bautista et
al., (2008).
Chelating Agents Degradation by Various Fenton Processes
There were a number of studies of advanced oxidation processes based on
Fenton oxidation to degrade or mineralize the chelating agents particularly EDTA.
Due to mineralizing ability of H2O2 for organic pollutants, H2O2 is considered as ecofriendly and safe reagent (Bautista et al., 2008).Without Fe2+ activation, excessive
concentration H2O2 in alkaline environment (pH=10) was unable to degrade 0.04 mM
EDTA. It was recommended that the use of an effective catalyst might increase the
conversion rate into more biodegradable decomposition products (Rämö & Sillanpää,
2001). However, with the presence of transition metals (Fe2+), treatment of waste
containing EDTA by chemical oxidation obtained 90% of EDTA was degraded at the
initial concentration of 70 mM in 45 min (Tucker et al., 1999).
A study on Fenton’s oxidation to degrade EDTA from bleaching wastewater
reported that an almost complete removal of EDTA was achieved at the H2O2
concentrations of 74 mM, the pH of 4, and the H2O2:Fe2+:EDTA ratio of 70:2:1
(Pirkanniemi et al., 2007). This result was comparatively higher than whose
previously accomplished by Tucker et al.(1999), indicating 90% of EDTA at an initial
concentration of 70 mM as provided in Table 2.2.Further study is needed to check the
applicability of this method for the treatment of real wastewater and to develop
heterogeneous catalysts for this process. In addition, conventional Fenton process has
been modified to Fenton-like, electro-Fenton and photo-Fenton processes by using
iron-supported catalyst like Fe(III) and zero-valent iron (ZVI) to improve efficiency
and sludge associated problem caused by conventional Fenton process (Neyens &
Baeyens, 2003; Bautista et al., 2008; Jiang et al., 2013; Zhou et al.,2009 & 2010).To
degrade 1 mM EDTA, oxygen activation scheme applied in zero-valent iron system
attained 95% of EDTA degradation at an initial concentration of 1 mM at pH 6.5
within 2.5 h (Noradoun & Cheang, 2005). In another study, Zhou et al.(2009) applied
an oxidative treatment by using heterogeneous ZVI and ultrasound to facilitate
reduction of O2 to H2O2. While being oxidized to Fe2+, ZVI induced series ofFenton-
20
like oxidation and degraded EDTA. In the system, EDTA acts as a complexing agent
with the dissolved Fe2+and generates H2O2. The result indicated that a lower EDTA
degradation (81%) at its concentration of 0.32 mM at pH 7.5 due to excessive iron
catalyst added in solution that prevented the formation of O-2-FeII/III EDTA, slowing
down EDTA degradation by Fenton-like oxidation.
The application of heterogeneous metallophthalocyanine (FePcS) in Fentonlike oxidation to degrade five different chelating agents including EDTA from
bleaching effluent was conducted. The rate of EDTA degradation was found to be
dependent on the concentration of Fe2+, H2O2, its molar ratio to the Fenton’s reagent,
pH, and temperature. Almost complete degradation of iron complexes of chelating
agents studied was remarkably obtained between 60% to 100% under pH 1.5 and
initial chelants concentration of 0.1M within a reaction time of 1 h. In addition, the
most relevant iron, manganese, sodium, copper and calcium EDTA complexes can be
successfully eliminated, the conversions being 93, 76, 68, 62 and 49%, respectively,
after 3h of reaction (Pirkanniemi et al., 2003). More description is detailed in table
2.2.
Application Fenton and Fenton-like reactions under UV-A irradiation to
degrade the 5 mM EDTA achieved 80% of EDTA removal with EDTA:Fe2+ and
EDTA:Fe3+ ratio of 1:1 with the initial peroxide concentration of 100 mM in 4 hours.
However, in both cases the reaction rates were increased after 4 hours irradiation with
the total EDTA mineralization of 92 % (Fe2+, Fe3+, Fe3++ Cu2+ system). The
photolysis of Fe(III)-EDTA complex in EDTA destruction can make use of high
peroxide concentration unnecessary. Photo-Fenton reaction was suitable for the
treatment of wastewater from cleaning and decontamination of nuclear power plant
because this wastewater contained small amount of Fe2+ and Fe3+ coming from
corrosion process (Ghiselli et al., 2004). For high iron content and organic citric acid
(8 % synthetic citric acid solution) in the derusting wastewater, UV photo-Fenton-like
oxidation was used because excessive amount of iron caused Fenton reaction occur
automatically when H2O2 was added. It was indicated that UV/H2O2/Fe3+ could
decomposed citric acid better than UV/H2O2 and Fe2+/H2O2. This is apparently due to
the important role of UV in allowing Fe3+ and H2O2 to function as strong oxidant in
producing radical chain reaction. In Fe2+/H2O2 system without UV, only 10% of
21
complex removed due chelating effects and precipitation. 93% COD reduction was
achieved for UV/H2O2/Fe3+ (Kim et al., 2010). Photo-Fenton oxidation with the
application of visible radiation, UV radiation, and sunlight achieved a complete
degradation of 20000 mg/L EDTA within 31, 6 and 3 hours, respectively. The kinetics
of photodegradation using solar-Fenton reaction follow the order of solar-Fenton >
UV (254 nm)- Fenton > Visible-Fenton. The pH changes from acidic to alkaline range
during the photo-Fenton process indicated loss of chelating ability of EDTA and
formation of amide was confirmed. Therefore, the design and treatment of large
volume of decontamination waste containing EDTA using a solar Fenton process is
easy, cost effective, and safe to operate (Chitra et al., 2011). Mechanism of UV
induced destruction, OH radical induced destruction, and ferric ion induced
destruction were implied for EDTA (Kim et al., 2010).
Metal chelating complexes are not be easily removed or degraded by a
single process. Therefore, a number of studies have incorporated Fenton reaction with
other treatment methods to improve its efficiency (Bautista et al., 2008). The
application of Fenton, Fenton-like, and advanced Fenton reactions followed by
hydroxide precipitation in removal of Ni from NiEDTA wastewater were conducted.
The complete disappearance of NiEDTA and 92% of Ni (II) removal were obtained.
Fenton and Fenton-like reactions were effective to degrade EDTA and the
fragmentation of NiEDTA freed up Ni(II) ion which was removed by precipitation.
Fenton-like process representing higher Ni(II) removal efficiency than Fenton process
can be attributed to the mechanism of ligand exchange. However, advanced Fenton
process (Fe0 + H2O2) shows higher removal efficiency of Ni (98.2%) and requires
lower H2O2 amount than Fenton or Fenton-like processes. COD decreased from 252
mg/L to 53.3 mg/L; indicating about 78.8% COD reduction. Lower percentage of
COD removal may be attributed to the formation of intermediates of acetate and
formate. Less than 0.03 mg/L of residue iron concentration was identified after Fenton
type processes, which required no further treatment options. This leads to a
conclusion that Fenton type processes seems to be an economically and
environmentally friendly process for remediation of strong stability chalated heavy
metal wastewater (Fu et al., 2009, 2012). The optimum operating parameters are also
provided in table 2.2.
22
Table 2.2 Summary of Fenton process for various EDTA complex wastewaters
Wastewater
Type
Pollutant
Concentration
Optimum Conditions
Efficiency
Reference
EDTA
70 mM
pH=4, T= 20 °C, [Fe2+]= 5 mM, [H2O2]= 100 mM,RT= 30
EDTA=90%
Tucker et al. (1999)
EDTA=90%
Pirkanniemi et
min,
Fe-EDTA
200 mM
pH=1.5, T= 40 °C, [Fe2+]= 0.03 mM, [H2O2]= 0.88 mM,
RT= 180 min
EDTA
76 mM
pH=3, T= 40 °C, [Fe2+]= 0.5 mM, [H2O2]= 18.5 mM, RT= 3
al.(2003)
EDTA=98%
min
EDTA
68.5 mM
pH=3, T= 40 °C, [Fe2+]= 0.04 mM, [H2O2]= 0.88 mM,
Pirkanniemi et
al.(2007)
EDTA=99%
Chitra et al. (2004)
RT= 720 min
EDTA
5 mM
pH=3, [Fe2+]= 200 mM, [H2O2]= 0.55 mM, RT= 240 min
EDTA=80%
Ghiselli et al.(2004)
Ni-EDTA
Ni=25 mg/L
pH=3, T= 40-50°C , [Fe2+/3+]= 1 mM, [H2O2]= 141 mM,
Ni=92%
Fu et al. (2009)
precipitation pH= 11, RT= 60 min
EDTA=100%
Ni=25 mg/L
pH=3, T= 40-50°C , [ZVI]= 2 g/L, [H2O2]= 35 mM,
Ni=98.2%
COD= 252 mg/L
precipitation pH= 11.5, RT= 60 min
COD=79%
Cu=225.3mg/L;
pH=2-5, T= 40-50°C , [Fe2+]/[H2O2] molar ratio = 2 ,
Cu=100%
Ni-EDTA
Cu-EDTA
COD=1096 mg/L [H2O2]:[COD]=0.2-0.3, RT= 60-80 min
COD=87%
Fu et al. (2012)
Lan et al.(2012)
24
The treatment of metal chelating complex wastewater is not only for metals
removal but also for organic compound degradation. Another study combined interior
microelectrolysis (IM) and Fenton oxidation-coagulation (IM-FOC) to treat EDTACu(II) containing wastewater. COD was used indirectly to determine the
concentration of EDTA species in the wastewater. IM process provide nearly
complete Cu(II) removal and yielded 336.1 mg/L Fe(II) concentration at very low pH
(pH=1.39) in accordance with IM reaction mechanism as reported in reviews (Ju et
al., 2011; Ju & Hu, 2011). The poor treatment performance of COD by IM, indicating
that EDTA species cannot be effectively decomposed into small biodegradable
organic molecules by IM process. The Fe(II)-rich effluent of IM was suitable for
direct treatment in a subsequent Fenton oxidation without Fe(II) addition or pH
adjustment. Under the optimal operating condition, Cu(II) and COD decrease from
225.3 mg/L and 1096.6 mg/L to 0 mg/L and 142.6 mg/L with overall removal
efficiency of 100% and 87%, respectively by IM-FOC process. After treatment, the
BOD5/COD ratio of wastewater was enhanced from 0 to 0.42, indicating that EDTA
was effectively oxidized in the combined system (Lan et al., 2012).
25
CHAPTER 3
RESEARCH METHODOLOGY
This chapter provides methodology, materials, and reagents required for this
study. Experimental variables were also determined. Experimental procedures,
analytical methods, and kinetic study were described as follows:
Derusting Wastewater Characteristics
The derusting wastewater used in this study was obtained from the Kation
Power Company, a cleaning service company, located in Rayong Province, Thailand.
This cleaning service company produces varying amount of wastewater according to
the numbers and types of cleaning processes. According to Huang et al. (2000), the
average cleaning wastewater is about 2300 m3 during each boiler cleaning. The
wastewater is originally produced from cleaning processes of pipes or boilers. The
wastewater taken from the company is stored temporarily in a storage tank for further
experiments. During the cleaning processes, various chemicals and chelating agent
(EDTA) are applied to remove rusts and to protect pipe and boiler from corrosion.
Furthermore, the derusting wastewater is in the dark red color due to high iron
content, which will form a complex with the EDTA.
Materials and Chemical Reagents
The reagents used in this study were the analytical grade reagents and used
without any further purification. Deionized or distilled water was used in all
experiments. Chemical reagents for Fenton and Fenton-like processes and chemical
reagents for wastewater parameters analysis were included as described and listed
below:
1.
Chemicals for Fenton and Fenton-like Processes
1.1 Hydrogen Peroxide (H2O2 -35% w/w),
1.2 Sodium Hydroxide (NaOH, 10N)
1.3 Sulfuric Acid (H2SO4, 5N)
1.4 Ferrous Sulfate (FeSO47H2O) for Fenton reaction
26
1.5 Manganese Dioxide (MnO2)
2. Chemicals for Parameters Analysis
2.1 COD
2.1.1 Standard Potassium Dichromate Digestion Solution
2.1.2 Sulfuric Acid reagent
2.1.3 Ferroin Indicator
2.1.4 Standard Ferrous Ammonium Sulfate (FAS) Titrant
2.2 Total Iron/Soluble Iron/Ferric/Ferrous Iron
2.2.1 Hydrochloric Acid (HCl) conc,
2.2.2 Hydroxylamine solution, AR Grade
2.2.3 Ammonium Acetate buffer solution
2.2.4 Sodium Acetate solution,AR Grade
2.2.5 Phenanthroline solution, AR Grade
2.2.6 Potassium Permanganate (KMnO4)
2.2.7 Stock Iron solution
3. Equipment and Materials
3.1
Jar Test apparatus (six paddles and six beakers with volume of 1L)
3.2
pH meter (EUTECH)
3.3
Multiparameter Photometer (Hana Instruments HI 83205-2008)
3.4
Analytical balance (OHAUS)
3.5
UV-Vis Spectrophotometer (Varian)
3.6
Turbidity meter (EUTECH)
3.7
Drying oven
3.8
Evaporating dishes
3.9
Suction flask
3.10 Desiccator
3.11 0.45m filter paper (GF/C )
3.12 Burette stand
3.13 Centrifugal machine (Harmonic Series)
3.14 Other glass wares (pipettes, burette, measuring cylinder,
volumetric flash, small beakers...)
27
Experimental Design and Procedure
Treatment efficiency of Fenton and Fenton-like reactions are the function of
the operating parameters including dosage of [H2O2], [Fe2+], initial pH, and reaction
time. Therefore, the variables of the experiment were classified and described as
follows:
a.
Independent Variables
-
Initial pH values: 2, 3, 4, 5, 6, 7
-
[Fe2+] concentrations: 0.005, 0.01, 0.05, 0.08, 0.1 and 0.15 M
-
[H2O2] concentrations 0.5, 1.0, 1.5, 2.0, 2.5, and 3.0 M
-
Precipitation pH values: 6, 7, 8, 9, 10, 11.
-
Reaction time: 20, 40, 60, 80, 100, and 120 min.
b. Dependent Variables
-
Total COD (TCOD), Soluble COD (SCOD), Total Iron, Soluble
Iron, Fe2+, Fe3+, Ammonium Nitrogen, Nitrite Nitrogen, Nitrate
Nitrogen, TSS, conductivity, and TDS as objective parameters
c.
Control Variables
-
Room temperature (28°C) corresponding to the wastewater
temperature during Fenton and Fenton-like processes.
-
Rapid mixing at 150 rpm for 2 min followed by slow mixing at 50
rpm.
-
Homogenous wastewater characteristic in all experiments.
The proposed experimental design was divided into 2 sets of experiments.
First set of experiment was referred as the Fenton-like reaction (addition of H2O2
only) by utilizing existing iron in the wastewater as catalyst. The second set of
experiment was referred as Fenton reaction (additions of both H2O2 and Fe2+). The
detail experimental design and procedures are provided as following:
Determine wastewater characteristics
For each experiment, the wastewater stored in the storage tank was
poured in a large tank and then mixed thoroughly so that the homogeneous
mixture was achieved. The sample was randomly collected for analyses of
28
wastewater characteristics. Various water quality parameters including total
COD, Soluble COD, total iron, soluble iron, Fe2+, Fe3+, ammonium nitrogen,
nitrite nitrogen, nitrate nitrogen, conductivity, TSS and TDS were
determined.
Hydroxide Precipitation of Iron Before Fenton and Fenton-like
Processes
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then started mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
2.
Adjusted the pH of wastewater with H2SO4 or NaOH to pH values
of 6, 7, 8, 9, 10, 11 in beaker No. 1, 2, 3, 4, 5, 6, respectively.
3.
Kept mixing the solution in each beaker at the mixing speed of 50
rpm for 15 min.
4.
At the end of mixing period, measured the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
5.
Stopped mixing and allowed the precipitates to settle for 30 minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
The brief experimental procedure for hydroxide precipitation before Fenton
and Fenton-like reactions is provided in figure 3.1.
29
500 mL of wastewater
Homogenueous Mixing at 50
rpm for 5 min
Mixing at 50 rpm for 15 min
pH adjustment to 6, 7, 8, 9,
10, 11
Parameter Analysis (COD,
Iron, Ammonia, Nitrite,
Nitrate)
Settling down for 30 min
Parameter Analysis (COD,
Iron, Ammonia, Nitrite,
Nitrate)
Figure 3.1 Hydroxide precipitation before Fenton and Fenton-like reactions
Effects of Initial pH on Fenton-like Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then started mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
2.
Adjusted the pH of wastewater with H2SO4 or NaOH to pH values
of 2, 4, 6, 8, 10, 12 in beaker No. 1, 2, 3, 4, 5, 6, respectively. Kept
mixing the solution in each beaker at the mixing speed of 50 rpm for
a few minutes. Then, measured the parameters such as pH, TDS,
conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
30
3.
Gradually added H2O2 at the concentration of 2.0 M into each
beaker. Kept mixing at the same speed for 60 minutes.
4.
At the end of mixing period, measured the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
5.
After collecting the samples, adjusted the pH to 8 to stop the
Fenton-like reaction. Continue mixing for another 15 minutes, and
then stop mixing and allowed the precipitates to settle for 30
minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of including TCOD, SCOD, total iron, soluble iron, Fe2+,
Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
7.
Repeated steps 1-7 with various pH values around the optimum pH
determined previously to obtain the best pH value.
Effects of H2O2 Concentration on Fenton-like Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank with H2SO4
or NaOH to the optimum pH value determined from the previous
study. Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
2.
Gradually added H2O2 with six different concentrations of 1.0, 1.5,
2.0, 2.5, 3.0, and 3.5 M into each beaker. Keeped mixing at the same
speed for 60 minutes.
8.
At the end of mixing period, measured the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
31
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
3.
After collecting the samples, adjusted the pH to 8 to stop the
Fenton-like reaction. Continued mixing for another 15 minutes, and
then stoped mixing and allowed the precipitates to settle for 30
minutes.
9.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen..
Effects of Reaction Time on Fenton-like Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank with H2SO4
or NaOH to the optimum pH value determined from the previous
study. Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
2.
Gradually added H2O2 at the optimum concentration determined
from previous study into each beaker. Kept mixing at the same
speed for 20, 40, 60, 80, 100, and 120 minutes of beaker No.1, 2, 3,
4, 5, and 6, respectively.
3.
After each mixing period of each beaker, measured the parameters
such as pH, TDS, conductivity in the beaker and then collect the
samples for additional analyses including TCOD, SCOD, total iron,
soluble iron, Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and
nitrate nitrogen
4.
After collecting the samples, adjusted the pH to 8 to stop the
Fenton-like reaction. Continued mixing for another 15 minutes, and
then stopped mixing and allowed the precipitates to settle for 30
minutes.
32
5.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
Hydroxide Precipitation of Iron After Fenton-like Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank to the
optimum pH value determined from the previous study and then
Filled every beaker with 500 mL of wastewater sample taken from
the large tank and then start mixing at 50 rpm for a few minutes to
have homogenous characteristic of wastewater
2.
Gradually added H2O2 at the optimum concentration determined
from previous study into each beaker. Keep mixing at the same
speed for a period of the optimum reaction time.
3.
After ending the mixing period, adjusted the pH of wastewater with
H2SO4 or NaOH to pH values of 6, 7, 8, 9, 10, 11 in beaker No. 1, 2,
3, 4, 5, 6, respectively.
4.
Kept mixing the solution in each beaker at the mixing speed of 50
rpm for 15 min.
5.
Stopped mixing and allowed the precipitates to settle for 30 minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
Effects of Initial pH on Fenton Process
1.
Prepareed a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
33
2.
Adjusted the pH of wastewater with H2SO4 or NaOH to pH values
of 2, 3, 4, 5, 6, 7 in beaker No. 1, 2, 3, 4, 5, 6, respectively. Kept
mixing the solution in each beaker at the mixing speed of 50 rpm for
a few minutes. Then, measureed the parameters such as pH, TDS,
conductivity in the beaker and then collect the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
3.
Adjusted the mixing speed to 150 rpm, and then added Fe2+ with a
concentration of 0.05 M. Maintained the mixing speed for 10
minutes to distribute the ferrous thoroughly in the beaker.
4.
After 10 minutes, adjusted the mixing speed to 50 rpm and
gradually added H2O2 at the concentration of 2.0 M into each
beaker. Kept mixing at the same speed for 60 minutes.
5.
At the end of mixing period, measured the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
6.
After collecting the samples, adjusted the pH to 8 to stop the Fenton
reaction. Continued mixing for another 15 minutes, and then
stopped mixing and allowed the precipitates to settle for 30 minutes.
7.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
8.
Repeated steps 1-8 with various pH values around the optimum pH
determined previously to obtain the best pH value.
Effects of Fe2+ Concentration on Fenton Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank with H2SO4
34
or NaOH to the optimum pH value determined from the previous
study. Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater
2.
Adjusted the mixing speed to 150 rpm, and then add Fe2+ with six
different concentrations of 0.005, 0.01, 0.05, 0.08, 0.1 and 0.15 M
into beaker No.1, 2, 3, 4, 5, and 6, respectively. Maintained the
mixing speed for 10 minutes to distribute the ferrous thoroughly in
the beaker.
3.
After 10 minutes, adjusted the mixing speed to 50 rpm and
gradually added H2O2 concentration of 2.0 M into each beaker.
Keep mixing at the same speed for 60 minutes.
4.
At the end of mixing period, measureed the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
5.
After collecting the samples, adjusted the pH to 8 to stop the Fenton
reaction. Continue mixing for another 15 minutes, and then stop
mixing and allow the precipitates to settle for 30 minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
Effects of H2O2 Concentration on Fenton Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank with H2SO4
or NaOH to the optimum pH value determined from the previous
study. Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater.
35
2.
Adjusted the mixing speed to 150 rpm, and then add Fe2+ with the
optimum concentration determined from previous study.
Maintained the mixing speed for 10 minutes to distribute the ferrous
thoroughly in the beaker.
3.
After 10 minutes, adjusted the mixing speed to 50 rpm and
gradually added H2O2 with six different concentrations of 1.0, 1.5,
2.0, 2.5, 3.0, and 3.5 M into each beaker. Kept mixing at the same
speed for 60 minutes.
4.
At the end of mixing period, measured the parameters such as pH,
TDS, conductivity in the beaker and then collected the samples for
additional analyses including TCOD, SCOD, total iron, soluble iron,
Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate
nitrogen.
5.
After collecting the samples, adjusted the pH to 8 to stop the Fenton
reaction. Continued mixing for another 15 minutes, and then
stopped mixing and allowed the precipitates to settle for 30 minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
Effects of Reaction Time on Fenton Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank with H2SO4
or NaOH to the optimum pH value determined from the previous
study. Filled every beaker with 500 mL of wastewater sample taken
from the large tank and then start mixing at 50 rpm for a few
minutes to have homogenous characteristic of wastewater.
2.
Adjusted the mixing speed to 150 rpm, and then add Fe2+ with the
optimum concentration determined from previous study. Maintain
the mixing speed for 10 minutes to distribute the ferrous thoroughly
in the beaker.
36
3.
After 10 minutes, adjusted the mixing speed to 50 rpm and
gradually added H2O2 at the optimum concentration determined
from previous study into each beaker. Kept mixing at the same
speed for 20, 40, 60, 80, 100, and 120 minutes of beaker No.1, 2, 3,
4, 5, and 6, respectively.
4.
After each mixing period of each beaker, measured the parameters
such as pH, TDS, conductivity in the beaker and then collected the
samples for additional analyses including TCOD, SCOD, total iron,
soluble iron, Fe2+, Fe3+, ammonium nitrogen, nitrite nitrogen, and
nitrate nitrogen.
5.
After collecting the samples, adjusted the pH to 8 to stop the Fenton
reaction. Continued mixing for another 15 minutes, and then
stopped mixing and allow the precipitates to settle for 30 minutes.
6.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
Hydroxide Precipitation of Iron After Fenton Process
1.
Prepared a Jar Test apparatus equipped with 6 beakers (size 1000-L
each). Adjusted the pH of wastewater in the large tank to the
optimum pH value determined from the previous study and then Fill
every beaker with 500 mL of wastewater sample taken from the
large tank and then started mixing at 50 rpm for a few minutes to
have homogenous characteristic of wastewater.
2.
Adjusted the mixing speed to 150 rpm, and then add Fe2+ with the
optimum concentration determined from previous study.
Maintained the mixing speed for 10 minutes to distribute the ferrous
thoroughly in the beaker.
3.
After 10 minutes, adjusted the mixing speed to 50 rpm and
gradually added H2O2 at the optimum concentration determined
37
from previous study into each beaker. Kept mixing at the same
speed for a period of optimum reaction time.
4.
After ending the mixing period, adjusted the pH of wastewater with
H2SO4 or NaOH to pH values of 6, 7, 8, 9, 10, 11 in beaker No. 1, 2,
3, 4, 5, 6, respectively.
5.
Kept mixing the solution in each beaker at the mixing speed of 50
rpm for 15 min.
6.
Stopped mixing and allow the precipitates to settle for 30 minutes.
7.
Collected the supernatant for sample analyses. The supernatant was
centrifuged at 2000 rpm and filtrated by 0.45m filter paper for
analyses of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen.
The experimental procedure of Fenton and Fenton-like reaction is
summarized in figure 2.3. It was noted that Fenton and Fenton-like reactions were
experimentally conducted in a similar procedure. However, the major difference was
that Fenton-like reaction ( add H2O2 only) was proceeded without adding Fe2+ during
the operation under mixing at 50 rpm, whereas Fenton reaction (H2O2 + Fe2+) required
to add Fe2+ under mixing at 150 rpm before addition of amount of H2O2.
38
500 mL of wastewater
into each beaker
Homogenueous Mixing
at 50 rpm for few
minute
Without adding
Fe2+/adding Fe2+ under
mixing at 150 rpm for
10 min
pH adjustment to 3
Addition of H2O2 under
mixing at 50 rpm for
60 min
Parameter Analysis
(COD, Iron, TDS, TSS,
Ammonia, Nitrite,
Nitrate)
Settling down for 30
min
pH adjustment to 8
Parameter Analysis
(COD, Iron, TDS,
TSS, Ammonia,
Nitrite, Nitrate)
Figure 3. 2 Experimental procedure for Fenton reaction (Adding Fe2+) and Fenton-like
reaction (Without adding Fe2+)
Optimum Conditions
The optimum condition was determined for each ferrous and hydrogen
peroxide concentration by computing the removal efficiencies of pollutants at
different varying concentration of reagents used in Fenton and Fenton-like oxidations.
The removal efficiency (R) is calculated by the following equation:
Removal Efficiency (R) =
A- B
100
A
39
where, A represents the initial characteristic of the objective parameters; B represents
the final characteristics of the objective parameters. The objective parameters include
the TCOD, SCOD, Total Iron, Soluble Iron, Fe2+, Fe3+, TSS, TDS, Ammonium
Nitrogen, Nitrite Nitrogen, and Nitrate Nitrogen.
Analytical Method
The analytical methods for each parameters were analyzed according to the
Standard Method for the Examination of Water and Wastewater (APHA, 2005). They
were briefly described as following:
1. The TCOD and SCOD of treated water was determined by the close reflux
titrimetric method (Method 5520).
2. The pH of solution was measured with a EUTECH pH meter.
3. Total iron, soluble iron, ferric and ferrous concentrations were analyzed by
Phenanthroline Method (Standard Method 3500).
4. Total suspended solid (TSS) was determined by standard method (Method
2540).
5. Total dissolved solid (TDS) was determined by portable TDS meter
(STARTER 300C)
6. Ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen were determined
by ion chromatography.
The detail description of each parameter analytical method was referred to
Appendix B.
Kinetic Study
In this study, practical reaction kinetic of Fenton and Fenton-like reactions
for organic compound degradation measured in COD reduction were determined from
the experimental effect of reaction time on both processes. The experiments were
conducted by varying reaction time of 20, 40, 60, 80 100 and 120 min as previously
discussed in order to monitor COD reduction and removal efficiency. As a result,
experimental data on COD reduction at different reaction time were observed.
40
The kinetics of Fenton and Fenton-like oxidations of COD removal
theoretically can be represented by the following nth-order reaction kinetics as
described in Skoog and West (2004), Bautista et al. (2007) and Wang (2008).
−
d(COD)
=k(COD)n
dt
where C represents the COD concentration, n is the order of the reaction, k is
the reaction rate coefficient and t is the time
The equation yields the following integrated equations respectively when it
is integrated for the zero, first, and second order.
(COD)t = -k t + (COD)0  (zero order reaction)
ln(COD)t = -kt + ln(COD)0  (first order reaction)
1
1
= kt +
 (second order reaction)
(COD)t
COD0
The obtained experimental data on COD concentration reduction with
respect to time were used to plot the curve and fit data points in Microsoft Excel,
2007 according to the zero order, first order and second order reaction as shown
above. The best fit is chosen when the coefficient of linearity (R2) is nearly equal to
the value of 1. Therefore, the kinetic rate constant (k) can be determined from the
slope of the linear line.
24
CHAPTER 4
RESULTS AND DISCUSSION
In this chapter, the results of the experiments are presented and discussed.
As stated in chapter 3, the wastewater characterization was conducted in order to
study the characteristic of wastewater and initial concentration of the parameters that
showed in the next section.
The second stage of the experiment was designed to investigate the possible
application of precipitation alone without combining with Fenton and Fenton-like
reactions (Precipitation before Fenton and Fenton-like reactions) for iron
precipitation. Therefore, the results of experiment will be shown in this chapter.
Third stage of the experiment was conducted to investigate the feasibility of
Fenton-like reaction by adding only H2O2 and utilizing the existing iron species in
wastewater as the catalyst to initiate the Fenton-like reaction. The impact of operating
parameters including initial pH, [H2O2], reaction time, and precipitation pH on
removal efficiency of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+, ammonium
nitrogen, nitrite nitrogen, and nitrate nitrogen will be discussed in detail in this
chapter.
Final stage of experiment was also conducted for Fenton reaction (additions
of both H2O2 and Fe2+). The results of the experiment on the impact of operating
parameters including initial pH, [Fe2+], [H2O2], reaction time, and precipitation pH on
removal efficiency of TCOD, SCOD, total iron, soluble Iron, Fe2+, Fe3+, ammonium
nitrogen, nitrite nitrogen, and nitrate nitrogen will also presented and discussed.
Finally, the result on the comparison of Fenton and Fenton-like reactions
will be discussed and the kinetic study of both processes are also conducted
accordingly.
Wastewater Characterization
The wastewater is originally produced from cleaning processes of boilers.
During the cleaning processes, the rust (iron oxide) is dissolved simultaneously by
using hydrochloric acid in the presence of chelating agent known as ethylenediamine
25
tetraacetic acid (EDTA) followed by application of ammonia to dissolve the metallic
copper and to avoid very low pH value of wastewater. The cleaning wastewater often
contains large amounts of iron and high concentration of chelating agents (Huang et
al., 2000; Bansal, 2012). Iron (Fe3+) is the most prevalent cation, generally present at a
concentration of 1000-10000 mg/L (Huang et al., 2000). According to analysis, the
average total iron concentration and TCOD of the sample are 3920 mg/L and 22257
mg/L, respectively. Ferric ion (Fe3+) reaches up to 3682 mg/L as provided in table 4.1.
All parameters value are shown to be higher than the limited effluent standard except
total suspended solid (TSS). Low TSS concentration indicated that derusting
wastewater are highly soluble and the pollutants are predominantly presented in the
dissolved form. High concentrations of iron are the most important characteristics of
the chemical cleaning wastewater. Kim et al, (2010) indicated that the derusting
wastewater is in the reddish dark color due to high iron concentration, which form a
complex with the chelating agent (EDTA). In addition, the pH value of wastewater is
about 10 due to the presence of the ammonia applied during the cleaning process
which results high concentration of ammonium (NH4+) up 16059 mg/L in the
wastewater. This high pH also indicates alkalinity which makes wastewater persistent
to low pH adjustment.
The derusting wastewater exceeds the permitted industrial effluent standard
(Thai effluent standard for industries: pH 6.5-8.5, COD < 400 mg/L, TSS < 150 mg/L,
TDS < 5000 mg/L and iron < 5 mg/L according to Pollution Control Department).
Therefore, this wastewater has to be treated before discharge into central wastewater
treatment facilities or natural water bodies in the environment.
Table 4-1 Derusting wastewater characteristics
Parameters
Average value
Limited effluent
pH
10.3
6.5-8.5
TCOD (mg/L)
22257
< 400
SCOD
22110
< 400
Total iron (mg/L)
3920
< 0.5
26
Dissolved iron (mg/L)
3801
0.5
Ferric (Fe3+) (mg/L)
3682
-
Ferric (Fe2+) (mg/L)
238
-
TDS (g/L)
9.84
< 5000
TSS (mg/L)
68
< 150
Conductivity (mS/cm)
19.69
-
Ammonium nitrogen (NH4+-N) (mg/L)
16059
< 1.1
Nitrite nitrogen (NH2- -N) (mg/L)
337
< 45
Nitrate nitrogen (NH3- -N) (mg/L)
208
-
Note: the limited effluent is based on Thai industrial effluent standard in Pollution
Control Department (PCD). www.pcd.go.th (Retrieved: September 20, 2013).
Hydroxide precipitation of iron before Fenton and Fenton-like
reactions
The presence of chelating agents, (EDTA and Citric acid) has commonly
known that conventional methods such as precipitation is inapplicable for metal
removal due to the complexing effect of the chelating agent toward heavy metals (Fu
et al., 2009; Chitra et al., 2012). Indeed, the derusting wastewater contains high
concentration of iron and chelating agent, EDTA, hydroxide precipitation may not
work for iron removal because of the Fe-EDTA complex as reported in literature (Fu
et al., 2012)
To investigate the performance of conventional hydroxide process in
removal of iron (Fe2+ /Fe3+ ) from derusting wastewater, the experiment was
conducted using hydroxide precipitation of iron before Fenton and Fenton-like
reactions by varying initial pH of wastewater at 6, 7, 8, 9 10, and 11 consecutively.
The objective of this experiment is not only to assess the hydroxide precipitation
ability in removal of iron but also to monitor the change of the initial concentration of
other parameters such as TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+,
27
ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen as the function of the
change in initial pH of the wastewater. The result is shown in table 4.2 below.
Table 4-2 The change of concentration in different precipitation pH
Precipitation pH
Parameters
pH=6
pH=7
pH=8
pH=9
pH=10 pH=11
TCOD (mg/L)
22500
24375
22500
22500
22500
24375
SCOD (mg/L)
24375
21750
21750
22500
21750
22500
Total Iron (mg/L)
3861
3932
3789
3968
3790
3665
Dissolved Iron (mg/L)
3623
3825
3706
3420
3754
3594
Fe2+ (mg/L)
1480
1206
1670
1170
1254
1301
Fe3+ (mg/L)
2381
2726
2119
2119
2537
2364
TDS (mg/L)
13.94
13.87
13.14
12.34
9.73
8.39
TSS (mg/L)
138
105.10
83.33
47.61
36.90
71.43
Conductivity (mS/cm)
27.85
27.7
26.3
24.7
19.47
16.79
NH4+ –N (mg/L)
13358
16130
17846
16157
17563
12324
NO2- –N (mg/L)
254
254
274
267
246
253
NO3-–N (mg/L)
29
39
47
22
28
19
It was indicated that the concentration of TCOD, SCOD, total iron and
dissolved iron remain almost unchanged after hydroxide precipitation before and after
precipitation process. However, the ferrous (Fe2+) concentration increased up to 1480,
1206, 1670, 1170, 1254, and 1301 mg/L at the precipitation pH of 6, 7, 8, 9, 10 and 11
respectively. This increase of ferrous ion concentration may be resulted from the
decrease of wastewater pH leading to the change of ferric ion to ferrous ion.
28
In addition, TDS and conductivity in the wastewater increase continuously
when pH of hydroxide precipitation was decreased or adjusted to lower pH. The
derusting wastewater initially contains 9.84 g/L of TDS and 19.69 mS/cm of
conductivity at the initial pH of 10. However, the decrease of pH from 10 to 6 in
precipitation process resulted in the increase of TDS and conductivity from 9.84 g/L
and 19.69 mS/cm to 13.94 mg/L and 27.85 mS/cm, respectively. This increasing
amount of TDS and electrical conductivity is apparently resulted from addition of acid
for pH adjustment (Deng & Englehardt, 2006)
The presence of EDTA cause metal complexation. The complexing ability of
the EDAT with iron is explained in the stability constant. The higher stability
constant, more complexing ability will be. The extent of complexation related to
metal-EDTA stability constant is provided as follow: log(K)Fe(III)-EDTA = 27.7 and
log(K)Fe(II)-EDTA = 16.0 (Skoog and West, 2004). Fe3+ has an EDTA log stability
constant considerably higher than Fe2+. In addition, Fu et al., (2009) also reported that
conventional precipitation process is not applicable to remove metal ions because
EDTA cause the dramatic increase in the solubility of heavy metal ion. The result
indicated that hydroxide precipitation is inapplicable or unable to remove the iron
from wastewater due to the complexing ability of the iron toward chelating agent,
EDTA.
Fenton-like reaction
As explained in chapter 3, Fenton-like reaction (addition of H2O2 only) was
experimentally conducted by utilizing existing iron (Fe3+ as predominant species) in
the wastewater as catalyst to react with H2O2 to generate hydroxyl radicals (•OH) for
organic complex degradation monitored in term of COD removal efficiency. The
major objective Fenton-like reaction is to evaluate the extent of treatment
performance efficiency in removal of TCOD, SCOD, total iron, soluble iron, Fe2+,
Fe3+, ammonium nitrogen, nitrite nitrogen, and nitrate nitrogen. The experimental
results and impact of operating parameters including initial pH, [H2O2], reaction time,
and precipitation pH was presented and discussed as following:
29
Effect of initial pH
pH is one of the major factors that limits the performance of Fenton-like
processes. It affects the speciation of iron and decomposition of hydrogen peroxide.
Efficiency of Fenton process is based on the pH and acidic pH highly favors the
oxidation reaction (Umar et al., 2010). To examine the effect of pH on Fenton-like
reaction on TCOD, SCOD, total iron and dissolved iron removal efficiencies, the
experiments were conducted at different initial pH of 2, 2.5, 3, 3.5, 4, 6, 8, 10 and 12
with 2 M H2O2. The results were indicated in figure 4-1 and figure 4-2 below.
As represented in figure 4-1, it is evident that Fenton-like reaction shows
high sensitivity with pH value and high reactivity at the pH range lower than 3.5. The
increase of removal efficiencies from 59.58 to 77.31% for TCOD and from 67.47 to
77.05 % for SCOD was obtained when pH increased from 2 to 3. The removal
efficiencies of TCOD and SCOD decreased from 77.32 to 43.96 % and from 77.31 to
43.59 %, respectively with the increase of pH from 3 to 4. At the pH greater than 6,
Fenton-like reaction did not achieve the TCOD and SCOD removal. It should be
noted that both TCOD and SCOD removal efficiencies by Fenton-reaction and
precipitation were similar indicating that Fenton-reaction was responsible for COD
removal particularly at pH 3.
In term of total iron and dissolved removal by Fenton-like reaction, the
results were indicated in figure 4-2 that both total iron and dissolved iron removal
efficiencies were 81.65% and 82.26%, respectively at initial pH of 3. The decreases of
removal efficiencies from 81.65 to 61.79 % for total iron and from 82.26 to 61.32 %
for dissolved iron were obtained when pH decreased from 3 to 2. However, the
increase of pH from 3 to 6 also resulted in decreasing the removal efficiencies of total
iron and dissolved iron from 81.65 to 27.03% and from 82.26 to 30.0%, respectively.
More importantly, further increase of the initial pH into alkaline condition resulted
poor removal efficiency (less than 15% of total and dissolved iron removal
efficiencies). This results is consistent with the literatures which reported that Fentonlike reaction could not work in alkaline pH range (Matthew Tarr, 2003; USperoxide,
2012). Figure 4-2 also indicated that less than 15% of total iron was used up and
removed by Fenton-like reaction alone and high concentration of iron in Fenton
30
treated effluent may be resulted from low solution pH leading to high solubility of
iron species (Fu et al., 2009; Lan et al., 2012).
Lower removal efficiency of TCOD, SCOD, total iron, and dissolved iron at
pH < 3 is due to stabilization of H2O2 as oxonium ions (H3O2+). The reaction between
•
OH and H+ also occurs as provide in reaction 4-1 and 4-2. Fe2+ regeneration by the
reaction of Fe3+ with H2O2 is inhibited at more acidic pH value (Wang et al., 2011).
H+
+ H2O2 →
H+
H3O2+
+ OH• → H2O
(4-1)
(4-2)
The decreases of removal efficiency at high pH (pH >6) due to selfdecomposition of H2O2 into oxygen and water at pH greater than 5, which reduces its
concentration in the solution (Bautista et al., 2008; Fu et al., 2009, 2012; Wang et al.,
2011). The results from this experiment are in agreement with those studies reported
by researchers, who found that acidic pH levels about 3 are usually optimum for
Fenton-like reaction (Babay et al., 2001; Fu et al., 2009).
31
(a)
(b)
Figure 4-1 Effect of initial pH on removal efficiencies of (a) TCOD and (b) SCOD at
[H2O2] of 2M, reaction time of 60 min, and precipitation pH 8
32
(a)
(b)
Figure 4-2 Effect of initial pH on removal efficiency of total iron and dissolved iron
by (a) Fenton-like reaction (b) precipitation at [H2O2] of 2M, reaction time
of 60 min, precipitation pH 8
33
Effect of H2O2 concentration
The amount of H2O2 is considered one of the most important factors in
Fenton-like reaction owing to its economic cost, sources of OH• generation. to
determine the suitable concentration of H2O2 required in Fenton-like reaction, the
experiments were conducted with various H2O2 of 0.5, 1, 1.5, 2, 2.5 and 3 M. The
results were shown in table 4-3 below.
The table 4-4 indicated the removal efficiencies of TCOD, SCOD, total iron,
and dissolved iron performed by Fenton-like reaction and by precipitation
subsequently. In Fenton reaction, it was clearly indicated that the removal efficiencies
of TCOD and SCOD increased as H2O2 concentration was increased. The increase of
H2O2 from 0.5 to 2.5M resulted in increasing of removal efficiencies from 51.15 to
83.94 % for TCOD and from 54.30 to 83.87% for SCOD, accordingly. Further
increase of H2O2 to 3M did not bring the removal efficiencies improvement for both
TCOD and SCOD. In addition, less than 21.26% of total iron and 33.7% of dissolved
iron were used up and removed by Fenton-like reaction. Therefore, iron concentration
always remained high in the Fenton-like effluent because narrow initial pH (pH 3)
operations cause high solubility of iron species as mentioned in previous work (Fu et
al., 2009).
Fenton-like treated effluent was immediately proceeded by subsequent
hydroxide precipitation. As presented in table 4-4, removal efficiencies of both TCOD
and SCOD were similar in both Fenton-like reaction and precipitation indicating COD
has already removed by Fenton-like reaction. However, the removal efficiencies of
total iron and dissolved iron were increased significantly after precipitation with the
increase of H2O2 concentration in Fenton-like reaction. For instance, the increase in
removal efficiencies from 44.66 to 87.30% for total iron and from 44.24 to 89.53%
for dissolved iron with the increase of H2O2 concentration from 0.5 to 2.5 M.
However, iron concentration remained unchanged with the further increase of H2O2.
In real application of Fenton-like reaction for complex industrial wastewater,
high concentration of H2O2 is required as the oxidizing agents and source of hydroxyl
radical (OH•) generation or hyperdroxyl radical (HO2•). Lower H2O2 applied could
result in inadequate radical generation, while overdosing could bring negative effects
34
on treatment performance (Bautista et al., 2008). In this study, lower removal
efficiencies were obtained when lower than 2.5 M H2O2 was applied due to the fact
that low OH• and HO2• radicals was generated. However, the application of H2O2
above 2.5 M did not enhance or improve the removal efficiency of TCOD, SCOD,
total iron and dissolved iron due to competitive reactions or scavenging effects of OH•
leading to negative impact on Fenton-like reaction (Neyens & Baeyens, 2003;
Matthew Tarr, 2003; Bautista et al., 2008; Wang, 2008). The competitive reactions
are shown below.
Fe2+ + OH• →
Fe3+ + OH−
k = 3.2 × 108 M-1s-1
(4-3)
H2O2 + OH• → HO2• + H2O
k = 2.7 × 107 M-1s-1
(4-4)
HO2• + OH• → O2
-
(4-5)
k = 5.2 × 109 M-1s-1
(4-6)
+ H2O
OH• + OH• → H2O2
It should be noted that Fenton-like reaction requires high amount of H2O2 in
the presence of EDTA in solution as previously reported by (Ghiselli, Jardim, Litter,
& Mansilla, 2004b; Fu et al., 2012). The optimum molar ratio of 40:1 in this study is
comparatively lower than those studies reported in literatures (Ghiselli et al., 2004; Fu
et al., 2009). Fu et al. (2009b) have also reported that the detrimental effect may be
observed when greater than 500:1 of H2O2: Fe2+ molar ratio is employed. However, in
this study lower ratio has been employed and detrimental effect has not been
observed.
For these low hydrogen peroxide concentrations, iron was detected as
ferrous iron once H 2 O 2 was completely consumed. This fact suggests that Fe 3+
ions, as well as other high-valent iron-oxo intermediates and ferryl complexes
probably reacted with the residual organic compounds present in the reaction medium
(Lunar et al., 2000).
35
(a)
(b)
Figure 4-3 Effects of H2O2 concentration on removal efficiencies of TCOD and
SCOD by (a) Fenton-like reaction (b) precipitation at initial pH of 3, precipitation pH
8
36
(a)
(b)
Figure 4-3 Effects of H2O2 concentration on removal efficiencies of total iron and
dissolved by (a) Fenton-like reaction (b) precipitation at initial pH of 3, precipitation
pH 8
37
Effect of reaction time
The reaction or oxidation time affects the removal efficiency in the treatment
of wastewater by Fenton process. If the oxidation time is too short, the organic
matters in wastewater would not be reacted with Fenton’s reagent completely, which
results in a bad pretreatment efficacy. By contraries, if the oxidation time is too long,
it is bound to an overlarge reactor and accordingly resulting in an increase of
investment, but the pretreatment efficacy does not improve significantly (Wu &
Wang, 2012). The optimum reaction time vary according to type and characteristic of
wastewater. Typical reaction times are 30-60 minutes for low strength wastewater.
For more complex or highly polluted industrial wastewater, the reaction may take
several hours. However, determination of reaction time still prove troublesome
(Matthew Tarr, 2003). In order to study the effect of reaction time for Fenton-like
reaction, a series of experiments were conducted by varying the reaction time from 20
min to 120 min and the results were shown in figure 4-4 below.
The results as presented in figure 4-5 clearly show that both TCOD and
SCOD removal efficiencies were increased with the increasing of the reaction time.
As increasing reaction time from 20 to 80 min, the removal efficiencies increased
from about 30 to 80 % for both TCOD and SCOD after Fenton-like reaction and
precipitation. Further extension of reaction time did not improve the removal
efficiencies. Similarly, the removal efficiencies of total iron and dissolved iron
increased as increasing the reaction time as presented in figure 4-5. Indeed, after
Fenton-like reaction approximated 50% of total iron and dissolved iron were removed
after 20 min onward and further explanation was shown in previous section. In
addition to Fenton-like reaction, subsequent hydroxide precipitation plays significant
role in removing the residue iron. As clearly indicated in figure 4-6, increasing the
reaction time from 20 to 80 min resulted in increasing removal efficiencies from
58.24 to 79.58% for total iron and from 57.37 to 80.35% for dissolved iron,
respectively.
38
(a)
(b)
Figure 4-3 Effect of reaction time (min) on removal efficiencies of (a) TCOD and (b)
SCOD at initial pH of 3, [H2O2] of 2.5 M and precipitation pH 8
39
(a)
(b)
Figure 4-4 Removal efficiencies (R %) of (a) total iron and (b) dissolved iron at
initial pH of 3, [H2O2] of 2.5 M, and precipitation pH 8
40
Effect of precipitation pH
pH adjustment subsequent to Fenton oxidation is typically conducted to
satisfy requirements for discharge, to stop the Fenton-like reaction and to convert
dissolved iron to iron sludge (Matthew Tarr, 2003; Neyens & Baeyens, 2003). The
experiments were conducted to study the effects of precipitation pH on the TCOD,
SCOD, total iron and dissolved iron removal efficiencies. In this experiment, Fentonlike treated effluent was immediately adjusted the pH of 6, 7, 8, 9, 10, and 11 by
adding sodium hydroxide to form precipitates. The results were indicated in figure 4-6
and 4-7 below.
The results presented in figure 4-5 (a) indicated that about 80 % removal
efficiencies of TCOD and SCOD were obtained by Fenton-like reaction. In addition,
Fenton-like effluents were subsequently proceeded by precipitation. Similar removal
efficiencies of TCOD and SCOD were obtained after precipitation indicating organic
degradation or total COD removal was accomplished by Fenton-like reaction before
precipitation process. On the other hand, about 48 % of total iron and as high as 58%
of dissolved iron were removed from the wastewater by the Fenton-like reaction as
presented in figure 4-6 (a). However, it is clearly indicated that both total iron and
dissolved iron removal efficiencies increased with the increase of precipitation pH in
precipitation process. As shown in figure 4-6 (b), the increase of precipitation pH
values from 6.0 to 11.0, the removal efficiencies of total iron increased from 49.16 %
to 95.44 % for total iron and from 49.54 % to 95.37% for dissolved iron after 80 min
of reaction time. At precipitation pH greater than 9, up to 93 % of both total iron and
dissolved iron removal efficiencies were achieved because the iron precipitates as
Fe(OH)3 at very high pH (Fu et al.,2009). Morgan & Lahav, (2007) also explained
that at the pH lower than 8.0, solubility and concentration of dissolved Fe2+ and Fe3+
remains high in solution. However, fraction of Fe2+ and Fe3+ are in solid phase or
form Fe(OH)2 and Fe(OH)3 as precipitates with high pH (pH>9).
It was concluded that about 95% of total iron and dissolved iron removal
efficiencies were attained at the precipitation pH of 11 for Fenton-like reaction. The
results from the experiment is constant with previous study of Fu et al., (2009) who
found the precipitation pH of 11 to be optimum for metal from complex wastewater.
41
(a)
(b)
Figure 4-5 Effect of precipitation pH on removal efficiencies (R %) of TCOD and
SCOD by (a) Fenton-like reaction (b) precipitation at initial pH of 3,
[H2O2] of 2.5 M, and precipitation pH 8
42
(a)
(b)
Figure 4-6 Removal efficiencies (R %) of total iron and dissolved iron by (a) Fentonlike reaction (b) precipitation at initial pH of 3, [H2O2] of 2.5 M, and
precipitation pH 8
43
Table 4-3 Concentration of total iron, Fe2+, and Fe3+ in Fenton-like effluent and
precipitation effluent at different precipitation pH
Precipitation
pH
Fenton-like effluent (mg/L)
Precipitation effluent (mg/L)
Total iron
Fe2+
Fe3+
Total iron
Fe2+
Fe3+
6
2030
675
1354
1993
884
1108
7
2066
627
1438
1886
606
1279
8
2108
657
1450
1033
624
409
9
2047
587
1459
650
575
75
10
2123
621
1501
256
212
44
11
2065
600
1464
178
175
3
It was also noted that the iron species including ferrous (Fe2+) and ferric
(Fe3+) were changed in the Fenton-like treated effluent and precipitation effluent.
Initially, total iron concentration existing in wastewater reaches up to 3920 mg/L. The
ferric (Fe3+) concentration is as high as 3682 mg/L which is the predominant species.
Therefore, ferrous (Fe2+) concentration is very low. However, ferrous concentration
increases continuously with the decrease of pH value of the wastewater. For instance,
decreasing the pH from 10 to 3 resulted in the increase of Fe2+ from 238 mg/L to 1609
mg/L. The increase of Fe2+ concentration after pH adjustment is due to the high
solubility of Fe2+ and the reduction of Fe3+ to Fe2+ become significant in narrow pH.
The results in table 4-4 indicated that about 48% of total iron was removed
and up to 2123 mg/L residue total iron was remained in Fenton-like treated effluent.
Among 2123 mg/L of total iron, the concentration of Fe2+ and Fe3+ were 621 mg/L
and 1501 mg/L indicating that Fe3+ is still be the predominant species in the Fentonlike treated effluent. Typically, Fe3+ predominates when the molar ratio of H2O2 to
total Fe is high; under those conditions reduction of iron by HO·2 (O−·2) is more
favorable. Regardless, HO·2(O−·2) radicals are decomposed to give one or the other
44
of the Fenton reactants, Fe(II) or H2O2, and thereby propagate the Fenton reaction
(Umar et al., 2010). In addition, Kim et al., (2010) also explained that the reduced
Fe2+ is quickly reoxidized to Fe3+ when the addition of H2O2 is presented which is
repeatedly occurring in term of a chain reaction as shown in reaction 2-6 and 2-6 in
the previous section.
TDS and Conductivity Content after Fenton-like reaction
One of drawback should be noted for wide application of Fenton treatment.
The significant quantities of acid are required to adjust the pH of wastewater resulted
in significant increase of the TDS and conductivity of the effluent leading to
operational hazards and safety and corrosion issues. TDS in effluent increased by
100% compared with that in influent in treating with Fenton method (Deng &
Englehardt, 2006). The effect of initial pH on concentration of TDS and conductivity
were indicated in figure 4-7.
Figure 4-7 Effects of initial pH on TDS and conductivity at [H2O2] of 2M, reaction
time of 60 min, and precipitation pH 8
45
The results apparently indicated that both TDS and electrical conductivity in
treated effluent increased continuously when initial pH was decreased or adjusted to
lower pH. The initial pH adjustment from 12 to 2 resulted in the increase of TDS from
11.96 g/L to 18.24 g/L and conductivity from 23.9 mS/cm to 38.5 mS/cm,
respectively. This increasing amount of TDS and electrical conductivity is apparently
resulted from addition of acid for pH adjustment and shown in reaction below (Deng
& Englehardt, 2006). The residue sulfate (SO42+) and sodium ion (Na+) resulted from
the use of pH adjusting agents remains in solution and contribute to increase of TDS
and electrical conductivity (Skoog and West, 2004).
H2SO4
NaOH
→
→
2H+
+
SO42+
(4-7)
OH−
+
Na+
(4-8)
Ammonia nitrogen, Nitrite, Nitrate removal
The application of Fenton process for ammonium, nitrite and nitrate
removals are limited in literatures. According to Huang et al. (2008), it was indicated
that ammonia could be oxidized by hydroxyl radical (OH•). A better removal of
ammonium was achieved and undetectable concentration of nitrite and nitrate was
observed (Lin & Chang, 2000). However, the increase of detected ammonium was
observed in Fenton treated effluent due to decomposition of organic nitrogen after 2
hour of start reaction and no nitrate was detected at any time. More importantly,
ammonium was found as intermediates of EDTA degradation by Fenton reaction,
while nitrite and nitrate was not detected in the solution (Babay et al., 2001; Sillanpää
et al., 2011).
In this study, the ammonium, nitrite and nitrate concentrations and removal
efficiencies were also determined.
The decline of ammonia concentration with Fenton and Fenton-like
reactions indicated that hydroxyl radical (OH•) could be employed to remove
ammonia. The results show that high pH was no doubt beneficial to ammonia
oxidation. It might be caused by the dissociative equilibrium of ammonia in water.
46
NH4+ + H2O
→ NH3 +
H3O+
pKa (NH4+)= 9.246
(4-9)
In light of the dissociation constant, the concentration of ammonia in
molecular form (NH3) and that of ammonia in the cation form (NH4+) were
approximately equal at pH 9.3. However, when the pH was 7.0, NH4+ was the
predominant component (99.5%) in the solution. Therefore, the amount of ammonia
in molecular form rose rapidly in basic condition. The acceleration of ammonia
removal with pH suggested that NH3 compared to NH4+ was more easily oxidized by
•OH.
Results show that the •OH could oxidize NH3 to NO2- and further to NO3.
Removal efficiencies of ammonia were low and were affected by initial pH value and
ammonia concentration. Results illustrate that •OH could oxidize NH3 to form •NH2.
•NH2, the main product of •OH with NH3, would further react with H2O2 to yield
•NHOH. Since •NHOH could not stay stable in solution, it would rapidly convert to
NH2O2-and consequently NO2-and NO3-.
OH•
→
•NH2
+
H2O
(4.9)
•NH2 + H2O2
→
•NHOH
+
H2O
(4.10)
NH3 +
•NH2 + OH•
→
NH2OH
(4.11)
When the organic substrate contains heteroatoms, mineralization often leads
to the formation of inorganic acids (HCl, HNO3, NH+4, H2SO4, etc.). Nitrogencontaining compounds may form HNO3exclusively (e.g., from ni- trophenols;
Kavitha and Planivelu, 2005) or a mixture of NH+4 and HNO3(e.g.from linuron, a
phenylurea herbicide; Katsumata et al., 2005). The influence of contaminant structure
and reaction conditions on the relative yields of NH+4and HNO3 are not well known.
Redox interconversion of NH+4 andNO−3 during HO·-intiated reactions involve a
number of intermediate steps and species (e.g., NH2OH, NH2, NO·) whose
importance is governed by pH, and presence of electron, proton, or hydrogen donors
47
or acceptors and O2 (Gonzalez et al., 2004). In the presence of O2 and absence of
organic matter, NH+4 is oxidized to NO−3, but this reaction can be very slow.
Sillanpää et al. (2011) reported that the by-products of EDTA oxidation such as NH4+,
NH3, NO3- causing incomplete oxidation of EDTA. The study of Chitra et al. (2011)
also confirmed that by stoichiometric calculation that, 17 mol of H2O2 are required to
completely oxidize 1 mol of EDTA to CO2, NH3, and H2O. Therefore, ammonia
(NH3) is also identified as the oxidation product as provide in reaction 4-1 below:
C10H14N2Na2O8 -2 H2O
+
H2O2
→
•NH2
+
H2O
(4-12)
Redox interconversion of NH+4 andNO−3 during HO·-intiated reactions
involve a number of intermediate steps and species (e.g., NH2OH, NH2, NO·) whose
importance is governed by pH. As nitrite accumulates, the oxidant acts as hydroxyl
radical scavenger leading to less hydroxyl radical generation and low removal
efficiency of organic compounds.
Denitrification using Fe2+ions as electron donors may be expressed as follows:
NO3-
+
5Fe2+
→
0.5N2
+
5FeOOH
+ 9H+
(4-1)
48
(a)
(b)
Figure 4-8 Removal efficiencies of ammonium and nitrate by (a) Fenton-like reaction
(b) precipitation at optimum condition
49
Kinetics of TCOD reduction by Fenton-like reaction
Due to the complexity of the organic pollutant in the derusting wastewater
and intermediates formed in the Fenton and Fenton-like reactions, it is impossible to
conduct a detailed kinetic study with the different individual reactions that take place
during the reaction. However, it is possible to conduct an approximated kinetic study
for organic compound degradation measured in COD removal (Wang et al., 2011;
Lucas & Peres, 2007). The complete oxidation reaction of Fenton and Fenton-like
reactions in removal of COD can be represented as below:
Organic matter (COD) + OH• → Oxidized product (P) + CO2 + H2O
(4-30)
In this study, kinetic analysis of Fenton-like and Fenton reactions were
conducted by monitoring the degradation of organic compound in term of TCOD
concentration reduction with time. Indeed, SCOD were not included because its
concentration reduction patterns were similar with TCOD.
In case of Fenton-like reaction, the kinetic analysis was carried out using the
experimental data on TCOD concentration reduction versus time obtained from
Fenton-like reaction performed at optimum initial pH of 3, and H2O2 of 2.5 M. In
addition, as discussed previously in the effect of reaction time on Fenton-like reaction,
TCOD concentration reduction and its removal efficiency improvement were
insignificant after 80 min of reaction. Therefore, optimum reaction of 80 min was
chosen for kinetic analysis. The experimental results on TCOD concentration
reduction and its removal efficiencies at different reaction time are indicated in table
4- below.
Table 4- Change of TCOD concentration with time after Fenton-like reaction
Time (min)
TCOD (mg/L)
Removal efficiency (%)
0
22257
0.00
20
15104
32.14
40
10746
51.72
50
60
5164
76.80
80
4299
80.69
The integrated forms of the zero, first, second order kinetic models as shown
equation were fitted to the experimental data to investigate the fitness of each kinetic
model to data points. The results on the kinetic model obtained are presented in table
4- . It is clearly indicated that the first-order kinetic model was fitted to the
experimental data with the good R-square value of 0.9732 as shown in figure 4- . The
zero-order kinetic model and the second-order kinetic model were fitted to the
experimental data with R-square value of 0.9310 and 0.9332, respectively. As a result,
this experiment shows that the rate of Fenton-like reaction is closer to the first-order
kinetic model.
(a)
51
(b)
(b)
Figure 4-7 ln (TCODt/TCOD0) versus time for first-order kinetic model at initial pH
of 3 and H2O2 of 2.5 M
52
Fenton reaction
As mentioned earlier that the experiment on Fenton reaction (additions of
both H2O2 and Fe2+) were conducted in order to evaluate its treatment performance in
removal efficiency of TCOD, SCOD, total iron, soluble iron, Fe2+, Fe3+, ammonium
nitrogen, nitrite nitrogen, and nitrate nitrogen. The experimental results and impact of
operating parameters including initial pH, [H2O2], reaction time, and precipitation pH
was presented and discussed as following:
Effect of Initial pH
Suitable initial pH plays a significant role in the treatment performance of
Fenton processes because it determines speciation of iron and decomposition of
hydrogen peroxide. The suitable pH for Fenton process is determined to be between 3
to 6 according to USperoxide (2012). However, acidic pH highly favors the Fenton
oxidation (Matthew Tarr, 2003; Neyens & Baeyens, 2003a). The oxidation potential
of hydroxyl radical decrease with increase in pH from Eo = 2.8 V to E14 = 1.96 V
(Umar et al., 2010).
To examine the effects of initial pH on TCOD, SCOD, total iron and
dissolved iron removal efficiencies, the experiments were conducted at different
initial pH of 2.0, 2.5, 3.0, 3.5 4.0, 4.5, 5.0, 6.0, and 7.0 with 2 M of H2O2 and 0.05 M
of Fe2+. It is evident in figure 4-9 (a) and 4-9 (b) that Fenton reaction provided high
reactivity and removal efficiencies of TCOD and SCOD at the pH ranged from 2.0 to
5.0. When pH increased from 2.0 to 3.0, the TCOD and SCOD removal efficiencies
increased from 83.54 to 91.77 % and from 83.43 to 91.71%, respectively. However,
the removal efficiency reductions were identified to be from 91.77 to 25.92% for
TCOD and from 91.71 to 28.59% with the increase of pH from 3.5 to 7.0 providing
the optimum initial pH of 3.0. Furthermore, as presented in figure 4-9 (a) and 4-9 (b),
removal efficiencies of TCOD and SCOD after Fenton reaction and after precipitation
process were similar indicating that Fenton reaction was responsible for organic
degradation monitored in term of COD reduction as reported in (Matthew Tarr, 2003).
More importantly, the removal efficiencies of total iron and dissolved iron
are different in both Fenton reaction and precipitation as indicated in figure 4-10 (a)
and (b). As much as 15% of total iron and dissolved was used up and removed by
53
Fenton reaction at initial pH ranged from 2 to 4.5 as presented in figure 4-10 (a).
Therefore, high concentration of iron in Fenton treated effluent was observed due to
low operating initial pH during Fenton reaction resulted in high solubility of iron
species in solution (Fu et al., 2009; Lan et al., 2012). In addition, the negative removal
efficiencies of both total iron and dissolved iron were attained at initial pH greater
than 5. Fenton reaction treated effluent then was subsequently proceeded by
hydroxide precipitation by adjusting pH to 8. The results in figure 4-10 (b) indicated
that when initial pH in the range from 2.0 to 4.0, the removal efficiencies of total iron
and dissolved were about 91% and 92%, respectively. Furthermore, at initial pH
greater than 5 resulted in negative removal efficiencies after both Fenton reaction and
precipitation. For instance, the pH ranged from 6.0-7.0, the residue total iron and
dissolved iron concentrations remained up to 4526 mg/L and 4222 mg/L, respectively.
Those residue irons were greater than initial total iron concentration (3920 mg/L) and
dissolved iron (3801 mg/L) leading to the negative removal efficiencies. Higher
amount of iron in solution was apparently resulted from the addition of 0.05 M Fe2+ as
the catalyst for Fenton reaction.
It is possibly to explain that lower removal efficiency of TCOD and SCOD
at pH less than 3.0 due to the following conditions. Extremely low pH values, the
[Fe(H2O)]2+ formed reacts relatively slowly with H2O2, producing less OH radical
(Umar et al., 2010). The stabilization of H2O2 as oxonium ions (H3O2+) and the
reaction between OH• and H+ occur as shown in reaction 4-1 and 4-2 in Fenton-like
reaction section (Wang et al., 2011). The removal efficiencies reduction at high pH
(pH > 6) due to Fe3+ precipitation as Fe(OH)3 and self-decomposition of H2O2 into
oxygen and water as shown in reaction (Bautista et al., 2008; Fu et al., 2009, 2012;
Wang et al., 2011). The results from this experiment are in agreement with those
studies reported by researchers (Neyens & Baeyens, 2003b; Fu et al., 2009b, 2012),
who found that acidic pH level about 3.0 is usually optimum for Fenton reaction.
H2O2 →
Fe3+ + OH-
H2O + 1/2 O2 (pH>5)
→
Fe(OH)3 (pH>6)
(4-14)
(4-15)
54
(a)
(b)
Figure 4-9 Effects of initial pH on removal efficiencies (R %) of TCOD and SCOD
by (a) Fenton reaction (b) precipitation at [Fe2+] of 0.05M, [H2O2] of 2 M,
and precipitation pH 8
55
(a)
(b)
Figure 4-10 Effects of initial pH on removal efficiencies (R %) of total iron and
dissolved iron by (a) Fenton reaction (b) precipitation at [Fe2+] of 0.05M,
[H2O2] of 2 M, and precipitation pH 8
56
Effect of Fe2+ concentration
Iron (Fe2+) is another important parameter in Fenton reaction that
catalytically decomposes hydrogen peroxide to generate hydroxyl radical (OH•).
Generally, removal of organics increases with increasing concentration of iron salt.
Inadequate concentration of iron in the operating condition will lead to insufficient
production of OH•, whereas overdosing of iron can favor the scavenging reaction
which prevents the reaction of OH• with contaminants resulting in poor treatment
efficiency (Matthew Tarr, 2003; Neyens & Baeyens, 2003).
To determine the effect of the Fe2+ concentration, the experiments were
conducted by varying the initial concentration of Fe2+ from 0.01 to 0.15M. The
removal efficiencies of TCOD, SCOD, total iron and dissolved iron after Fenton
reaction and precipitation are presented in figure 4-12 and 4-13 below.
The result apparently indicates that TCOD and SCOD removal efficiencies
increased with the increase of Fe2+ concentration. As presented in figure 4-12 (a) and
(b), the increase of Fe2+ from 0.01 to 0.05 resulted in the increase of removal
efficiencies from 72.82 to 91.39 % for TCOD and from 80.43 to 91.16% for SCOD,
respectively after Fenton reaction and precipitation. It should be noted that similar
removal efficiencies between TCOD and SCOD showed that most of organic
compound was in the dissolved form. However, lower removal efficiencies TCOD
and SCOD (R% < 91%) were observed with further increase of initial Fe2+
concentration.
It is generally accepted that the addition of Fe2+ as the catalyst for Fenton
reaction resulted in the accumulation of residue soluble iron in Fenton reaction and
amount of iron sludge during precipitation process (Bautista et al., 2008). The result
in figure 4-13(a) clearly indicated that both total iron and dissolved iron concentration
increased continuously as the increase of initial Fe2+ in Fenton reaction treated
effluent. In addition, negative removal efficiencies of total iron and dissolved iron
were observed when greater than 0.08M of Fe2+ was added in Fenton reaction.
Specifically, as high as 6645 mg/L of total iron and 5192 mg/L of dissolved iron were
detected in Fenton treated effluent which is comparatively higher than initial total iron
(3920 mg/L) and dissolved iron (3801 mg/L). However, the removal efficiencies of
total iron and dissolved iron increased after subsequent precipitation as presented in
57
figure 4-13(b). The results clearly indicated that the increase the initial Fe2+
concentration from 0.01 to 0.05 M resulted in increase of removal efficiencies from
66.61 to 90.46% for total iron and from 69.24 to 90.72% for dissolved iron,
accordingly. Further increasing of initial Fe2+ did not bring the improvement in
removal efficiencies of both irons; otherwise, lower removal efficiency was identified.
It was also observed in figure 4-1 (a) and (b) that when greater than 0.08
mg/L of initial Fe2+ concentration was applied the differences in removal efficiency
of total iron and dissolved iron were observed. This indicated that larger amount of
initial Fe2+ used during the Fenton reaction led to accumulation of residue colloidal
iron in precipitation effluent as stated in literature (Neyens & Baeyens, 2003; Fu et
al., 2009; USperoxide, 2012).
In this study, lower removal efficiencies of COD and iron were obtained
when lower than 0.05M of initial Fe2+ was applied. This apparently resulted from
lower generation of OH• and HO2• radicals because Fe2+ concentration were not
sufficiently used as catalyst for decomposition of H2O2 (Matthew Tarr, 2003). On the
other hand, lower removal efficiencies of TCOD, SCOD, total iron and dissolved iron
were attained due to the self-scavenging of OH• and HO2• radicals by Fe2+ and Fe3+.
The scavenging effects are similar in case of H2O2 overdosing leading to unnecessary
OH• and HO2• consumption (Bautista et al., 2007, 2008; Wang, 2008; Lou & Huang,
2009; Fu et al.,2009, 2012; Wang et al., 2011). The radical scavenging reactions are
shown in reaction 4-16 below.
Fe2+ + •OH → Fe3+ + OH−
k=3.2 × 108 M-1s-1
(4-16)
Fe2+ + HO2• → Fe3+ + HO2−
k= 1.2 × 106 M-1s-1
(4-17)
k= 2 × 103 M-1s-1
(4-18)
Fe3+ + HO2• → Fe2+ + H+ + O2
In addition, excess Fe2+ concentration contributes to an increase in effluent
TDS and electrical conductivity, as well as in the amount of iron sludge that requires
further treatment (Umar et al., 2010). The results clearly showed that increasing Fe2+
from 0.01 to 0.15 M resulted in the increase of TDS concentration from 16.56 to
17.71 g/L and electrical conductivity from 33.1 to 35.6 mS/cm ( Figure was not
58
attached). This increase of both TDS and electrical conductivity resulted from the use
of FeSO4 as the source of Fe2+ reagent (Umar et al., 2010; USperoxide, 2012).
Therefore, dissolution of FeSO4 remained the anion SO42- in solution. SO42- ion
physically contributes the increase of TDS and electrical conductivity in treated
effluent (USperoxide, 2012).
59
(a)
(b)
Figure 4-11 Effects of initial Fe2+ concentration on removal efficiencies (R %) of
TCOD and SCOD by (a) Fenton reaction (b) precipitation at initial pH of
3, [H2O2] of 2 M, and precipitation pH 8
60
(a)
(b)
Figure 4-12 Effects of initial Fe2+ concentration on removal efficiencies (R %) of
total iron and dissolved iron by (a) Fenton reaction (b) precipitation at
initial pH of 3, [H2O2] of 2 M, and precipitation pH 8
61
Effect of H2O2: Fe2+ molar ratio
In the Fenton process, the molar ratio of H2O2:Fe2+ is very important in
terms of overall cost and removal efficiency of the process. Excessive or shortage of
any of these two reagents results in the occurrence of scavenging reactions. It is
difficult to specify the optimal H2O2: Fe2+ molar ratio for wastewater treatment;
however, It varied in term of the degradation of different pollutant covering the range
from 1:1 to 1:400 (Wang et al., 2011). In this study, various molar ratios of H2O2:
Fe2+ were applied to the industrial wastewaters with the optimum initial pH of 3.0 and
0.05 M of initial Fe2+ determined from previous experiment. The effects of H2O2:Fe2+
on the removal efficiencies of TCOD, SCOD, total iron and dissolved iron are shown
in figure 4-14 and 4-15. The results indicate that the removal efficiencies of TCOD
and SCOD increased with the increase of the H2O2: Fe2+ molar ratios. As explained
previous section, the TCOD and SCOD removal efficiencies were obtained similarly
after Fenton reaction and precipitation. The increasing the H2O2: Fe2+ molar ratio
from 10 to 40 resulted in removal efficiency improvement of TCOD from 47.82 to
93.04 % and SCOD from 52.73 to 93 %, respectively. However, further increase of
H2O2:Fe2+ molar ratios above 40 did not enhance the removal efficiencies of both
TCOD and SCOD due to scavenging effects of hydroxyl radical by Fe2+ according to
reaction 4-16 as reported in literature reaction (Neyens & Baeyens, 2003; Matthew
Tarr, 2003; Bautista et al., 2007 and 2008; Wang, 2008).
As mentioned earlier, as high as 15% of both irons were used up and
removed by Fenton reaction. As indicated in figure 4-15 (a), less than 20% of total
iron and 33% of dissolved iron removal efficiencies was observed after Fenton
reaction. These removal efficiencies of both irons are resulted from the formation of
iron species solid particles and precipitates (Umar et al., 2010) .Therefore, lower
extent of both iron removal efficiencies after Fenton reaction were determined and the
difference between total iron and dissolved iron were obvious. In comparison to
Fenton reaction, higher removal efficiencies of total iron and dissolved iron were
obtained due the significant function of precipitation for iron removal efficiency as
clearly explained in the work of Fu et al., (2009; 2012). As seen in figure 4-15 (b), it
is evident that the improvements of removal efficiencies from 11.95 to 90.46% for
total iron and from 12.89 to 92.89% for dissolved iron were obtained with the increase
62
of H2O2: Fe2+ molar ratio from 10 to 40, respectively. On the other hand, almost
constant removal efficiencies of total iron and dissolved iron were achieved when
H2O2: Fe2+ molar ratio were greater than 40. Insignificant improvement or almost
constant removal efficiencies are apparently due to competitive reactions or
scavenging effects as previously shown and discussed in reaction (4-16). In addition,
the pH drop during Fenton reaction was evident as discussed previously. When the
H2O2: Fe2+ molar ratio was greater than 40, the increase of pH was observed due to
scavenging reaction of OH• by Fe2+ yields OH− as the product.
It should be noted that Fenton reaction requires high amount of H2O2 in the
presence of EDTA in solution as previously reported by Ghiselli et al. (2004) and Fu
et al. (2012). The optimum molar ratio of 40:1 in this study is comparatively lower
than those studies reported in literatures (Ghiselli et al., 2004; Fu et al., 2009). Fu et
al. (2009b) have also reported that the detrimental effects may be observed when
greater than 500:1 of H2O2: Fe2+ molar ratio is employed. However, in this study
lower ratios were employed and detrimental effects has not been observed.
63
(a)
(b)
Figure 4-13 Effects of H2O2:Fe2+ molar ratio on removal efficiencies (R %) of TCOD
and SCOD at initial pH of 3, [Fe2+] of 0.05M, and precipitation pH 8
64
(a)
(b)
Figure 4-14 Effects of H2O2:Fe2+ molar ratio on removal efficiencies (R %) of total
iron and dissolved iron by (a) Fenton reaction (b) precipitation at initial
pH of 3, [Fe2+] of 0.05 M, and precipitation pH 8
65
Effect of Reaction Time
The reaction time for a completion of Fenton reaction depends on the its
reagents (Fe2+ and H2O2) because the contaminant degradation rate is proportional to
the hydroxyl radical produced (Matthew Tarr, 2003). Even though Fenton reaction is
commonly known as fast reaction in term of organic degradation, determination of
reaction time for Fenton reaction found differently in literatures (Matthew Tarr, 2003
Bautista et al., 2008). In this study, the experiments were conducted by varying the
reaction time from 20 min to 120 min and the results were shown in figure 4-16 and
figure 4-17 below.
It is evident that the about 91% of TCOD and SCOD removal efficiencies
were achieved at the first 20 min of Fenton reaction while prolonging the reaction
time did not improve the removal efficiencies of both TCOD and SCOD as provide in
figure 4-16(a). The TCOD and SCOD removal efficiencies were obtained similarly
after Fenton reaction and precipitation as provided in figure 4-16(a) and 4-16(b).
Similarly, above 91% of the total iron and dissolved iron removal efficiencies were
also obtained in first 20 min of reaction after precipitation as shown in figure 4-17 (b).
In this study, the reaction time for Fenton reaction was found to be shorter than those
studies reported by researchers (Pirkanniemi et al., 2003; Fu et al., 2009; 2012; Lan et
al., 2012) who found that the required reaction time of longer than 60 min was
required for EDTA-complex wastewater treatment. However, the results in this study
are in agreement with the studies of San Sebastián Martinez et al.(2003) and Jiang et
al. (2013) who achieved optimum efficiencies in less than 20 min of Fenton reaction
due to the fast reaction in the first stage of Fenton oxidation.
66
(a)
(b)
Figure 4-15 Effects of reaction time on removal efficiencies (R %) of TCOD and
SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of
0.05 M, H2O2:Fe2+ molar ratio of 40, and precipitation pH 8
67
(a)
(b)
Figure 4-16 Effects of reaction time on removal efficiencies (R %) of total iron and
dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of 3,
[Fe2+] of 0.05 M, H2O2:Fe2+ molar ratio of 40, and precipitation pH 8
68
Effect of precipitation pH
The subsequent hydroxide precipitation is applied in order to precipitate and
remove iron from the Fenton treated effluent by adjusting solution pH to alkaline
range (pH >8). As discussed previously, solid precipitation as Fe(OH)3 was done by
hydroxide precipitation; however, solid formation and settling depend on precipitation
pH according to Fu et al., (2009). In this study, the experiments were conducted to
study the effect of precipitation pH on the removal efficiency of total iron, dissolved
iron, TCOD and SCOD.
The results indicate that about 15% of total iron and dissolved iron were
used up and removed from Fenton treated effluent as shown in figure 4-19 (a). High
concentration of iron in Fenton treated effluent may be resulted from low solution pH
leading to high solubility of iron species (Fu et al., 2009; Lan et al., 2012). However,
as can be seen in figure 4-19 (b) with precipitation pH value increasing from 6 to 9,
the removal efficiencies of total iron and dissolved iron increased from 55 % to 94 %
from after 20 min of reaction. When pH increased further, the removal efficiency of
total iron and dissolved iron was insignificantly improved because at very high pH,
the iron precipitates as Fe(OH)3 (Fu et al.,2009 ;2012). This is apparently explained
that at the pH lower than 8, solubility and concentration of dissolved Fe2+ and Fe3+
remain higher in solution. However, when pH of solution is greater than 8, fraction of
Fe2+ and Fe3+ are in solid phase or form Fe(OH)2 and Fe(OH)3 as precipitates as
shown in reaction (4-19)-(4-21) below (Morgan & Lahav, 2007; Fu et al., 2009b). On
the other hand, removal efficiency of TCOD and SCOD remained almost unchanged
(R = 90%) after Fenton reaction and precipitation in all pH values. Organic
compounds degradation or COD removal was already done by Fenton reaction before
precipitation process as explained previously.
Fe2+ +
2OH−
→
Fe(OH)2↓
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3↓
Fe3+
+
3OH− →
Fe(OH)3↓
(4-19)
(4-20)
(4-21)
In summary, precipitation does not improve the removal efficiency in terms
of total COD reduction; however, it is responsible for iron precipitation found in this
69
study. As reported, Fenton reaction is employed for organic degradation while the
alkaline range of pH is to stop Fenton reaction and precipitate irons (Lan et al., 2012;
USperoxide, 2012).
In addition to total iron determination, ferrous (Fe2+) and ferric (Fe3+)
concentrations were also measured in order to study the interconversion or
transformation among these two ions during Fenton reaction and precipitation. As it
was indicated that the initial concentration total iron is 3920 mg/L and as high as 3682
mg/L is in the form of Fe3+ which is predominant in the derusting wastewater. The
results of total iron, ferrous (Fe2+) and ferric (Fe3+) concentrations after Fenton
reaction and precipitation are indicated in figure 4-20 below. At the optimum
condition for Fenton reaction was adopted and conducted. As seen in figure 4-20 (a),
Fenton reaction achieved about 15% of total iron removal efficiency indicating about
3309 mg/L of total iron remained in Fenton treated effluent. Among 3309 mg/L of
total iron, about 2145 mg/L was in the form of Fe3+, while the remaining iron was
Fe2+ form. The increase of Fe2+ concentration was identified but its proportion was
about 2 times lower than Fe3+ concentration after Fenton reaction. The increase of
Fe2+ proportion in Fenton treated effluent may resulted from lower initial pH operated
during Fenton reaction (Morgan & Lahav, 2007). It is also resulted from addition of
Fe2+ as the catalyst for Fenton reaction and the reducing activity from Fe2+ to Fe3+ is
inhibited at low pH (Bautista et al., 2008). Moreover, Bautista et al. (2007) explained
that a high concentration of organic matter favors the regeneration of Fe2+ from the
reaction of Fe3+ and organic radicals as shown in reaction (2-3) in chapter 2. However,
larger proportion of Fe3+ concentration after Fenton reaction is resulted from molar
high ratio of H2O2 to total Fe2+ ; under those conditions reduction of iron by HO2• is
more favorable (Umar et al., 2010). Lunar et al.(2000) found that Fe2+ concentration
was not detectable in the course of Fenton reaction. This behavior, in the excess of
H2O2, is according to Fenton-like mechanism reactions. In addition, Kim et al., (2010)
also explained that the reduced Fe2+ is quickly reoxidized to Fe3+ when the addition of
H2O2 is presented which is repeatedly occurring in term of a chain reaction as shown
in reaction (2-6) and (2-8) in the previous section.
However, total iron, ferrous (Fe2+) and ferric (Fe3+) concentrations were
reduced continuously with the increase of precipitation pH as provided in figure 4-20
70
(b). The reduction of total iron concentration from 1763 to 219 mg/L was determined
when the precipitation pH was increased from 6 to 9. At the precipitation lower than
8, the proportion of Fe3+ concentration is still higher than Fe2+. When precipitation pH
was greater than 8, proportion of Fe3+ concentration was lower than Fe2+. At pH
solution greater than 8, fraction of Fe2+ and Fe3+ are in solid phase or form Fe(OH)2
and Fe(OH)3 as precipitates as indicated in reaction (4-19)-(4-21). However, Fe3+ are
easily precipitate as Fe(OH)3 than Fe2+ does when pH is greater than 8. This indicates
that Fe3+ was more subjective to precipitation than Fe2+ at the precipitation greater
than 8 (Morgan & Lahav, 2007)
71
(a)
(b)
Figure 4-17 Effects of precipitation pH on removal efficiencies (R %) of TCOD and
SCOD by (a) Fenton reaction (b) precipitation at initial pH of 3, [Fe2+] of
0.05 M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20 min
72
(a)
(b)
Figure 4-18 Effects of precipitation pH on removal efficiencies (R %) of total iron
and dissolved iron by (a) Fenton reaction (b) precipitation at initial pH of
3, [Fe2+] of 0.05 M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20
min
73
(a)
(b)
Figure 4-19 Concentration of total iron, Fe2+, and Fe3+ in (a) Fenton reaction effluent
(b) precipitation effluent at different precipitation pH at initial pH of 3,
[Fe2+] of 0.05 M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20
min
74
TDS and Conductivity Content after Fenton reaction
The increase of TDS and electrical conductivity is considered to be another
limitation for Fenton reaction in addition the sludge accumulation problem which
further treatment is requires due to operational hazards and safety and corrosion
issues. The pH adjustment agents including H2SO4 and NaOH are the major
contributions for the increase of TDS concentration and conductivity (Deng &
Englehardt, 2006; Umar et al., 2010). The TDS concentration and conductivity after
Fenton reaction and precipitation at different initial pH are indicated in figure 4-21
below.
Figure 4-20 Effects of initial pH on TDS and conductivity contents at [H2O2] of 2M,
reaction time of 60 min, and precipitation pH 8
The results apparently indicated that both TDS and electrical conductivity
increased continuously with the decreasing of initial pH during Fenton reaction and
increased further after precipitation. The initial pH adjustment from 7 to 2 resulted in
the increase of TDS from 9.94 g/L to 18.36 g/L and conductivity from 19.26 mS/cm
to 36.7 mS/cm, respectively after Fenton reaction. Further increase of TDS
concentration from 11.96 to 19.24 g/L and conductivity from 23.9 to 38.5 were
75
observed after precipitation. This increasing amount of TDS and electrical
conductivity is apparently resulted from addition of acid for pH adjustment and FeSO4
as the source of Fe2+ during Fenton reaction (Deng & Englehardt, 2006). The further
increase of TDS and conductivity identified after precipitation is clearly resulted from
addition of NaOH during precipitation process. The residue sulfate (SO42+) and
sodium ion (Na+) resulted from the use of pH adjusting agents remains in solution and
contribute to increase of TDS and electrical conductivity as shown in reaction 4-7 and
4-8 (Skoog and West, 2004).
Ammonium nitrogen, Nitrite, Nitrate removal
The decline of ammonia concentration with Fenton and Fenton-like
reactions indicated that hydroxyl radical (OH•) could be employed to remove
ammonia. The results show that high pH was no doubt beneficial to ammonia
oxidation. It might be caused by the dissociative equilibrium of ammonia in water.
NH4+ + H2O
→ NH3 +
H3O+
pKa (NH4+)= 9.246
(4-1)
In light of the dissociation constant, the concentration of ammonia in
molecular form (NH3) and that of ammonia in the cation form (NH4+) were
approximately equal at pH 9.3. However, when the pH was 7.0, NH4+ was the
predominant component (99.5%) in the solution. Therefore, the amount of ammonia
in molecular form rose rapidly in basic condition. The acceleration of ammonia
removal with pH suggested that NH3 compared to NH4+ was more easily oxidized by
•OH.
Results show that the •OH could oxidize NH3 to NO2- and further to NO3.
Removal efficiencies of ammonia were low and were affected by initial pH value and
ammonia concentration. Results illustrate that •OH could oxidize NH3 to form •NH2.
•NH2, the main product of •OH with NH3, would further react with H2O2 to yield
•NHOH. Since •NHOH could not stay stable in solution, it would rapidly convert to
NH2O2-and consequently NO2-and NO3-.
76
OH•
→
•NH2
+
H2O
•NH2 + H2O2
→
•NHOH
+
H2O
NH3 +
•NH2 + OH•
→
(4.1)
NH2OH
When the organic substrate contains heteroatoms, mineralization often leads to the
formation of inorganic acids (HCl, HNO3, NH+4, H2SO4, etc.). Nitrogen-containing
compounds may form HNO3exclusively (e.g., from ni- trophenols; Kavitha and
Planivelu, 2005) or a mixture of NH+4 and HNO3(e.g.from linuron, a phenylurea
herbicide; Katsumata et al., 2005). The influence of contaminant structure and
reaction conditions on the relative yields of NH+4and HNO3 are not well known.
Redox interconversion of NH+4 andNO−3 during HO·-intiated reactions involve a
number of intermediate steps and species (e.g., NH2OH, NH2, NO·) whose
importance is governed by pH, and presence of electron, proton, or hydrogen donors
or acceptors and O2 (Gonzalez et al., 2004). In the presence of O2and absence of
organic matter, NH+4 is oxidized to NO−3, but this reaction can be very slow.
Sillanpää et al. (2011) reported that the by-products of EDTA oxidation such as NH4+,
NH3, NO3- causing incomplete oxidation of EDTA. The study of Chitra et al. (2011)
also confirmed that by stoichiometric calculation that, 17 mol of H2O2 are required to
completely oxidize 1 mol of EDTA to CO2, NH3, and H2O. Therefore, ammonia
(NH3) is also identified as the oxidation product as provide in reaction 4-1 below:
C10H14N2Na2O8 -2 H2O
+
H2O2
→
•NH2
+
H2O
(4-1)
Redox interconversion of NH+4 andNO−3 during HO·-intiated reactions
involve a number of intermediate steps and species (e.g., NH2OH, NH2, NO·) whose
importance is governed by pH. As nitrite accumulates, the oxidant acts as hydroxyl
radical scavenger leading to less hydroxyl radical generation and low removal
efficiency of organic compounds.
Denitrification using Fe2+ions as electron donors may be expressed as follows:
77
NO3-
+
5Fe2+
→
0.5N2
+
5FeOOH
+ 9H+
(4-1)
According to (Lin & Chang, 2000), better removal of ammonium was
achieved and undetectable concentration of nitrite and nitrate was observed. However,
Lunar et al. (2000) detected the ammonium in Fenton treated effluent due to
decomposition of organic nitrogen after 2 hour of start reaction and no nitrate was
detected at any time. More importantly, ammonium was found as intermediates of
EDTA degradation by Fenton reaction, while nitrite and nitrate was not detected in
the solution (Babay et al., 2001).
78
(a)
(b)
Figure 4-21 Removal efficiencies of ammonium and nitrate by (a) Fenton reaction (b)
precipitation at optimum condition pH at initial pH of 3, [Fe2+] of 0.05
M, H2O2:Fe2+ molar ratio of 40 and reaction time of 20 min
79
Kinetics of TCOD reduction by Fenton reaction
In comparison to Fenton-like reaction, the kinetic analysis of Fenton
reaction was also conducted. As stated in chapter 3, the experiments were conducted
by varying reaction time of 20, 40, 60, 80 100 and 120 min as previously discussed in
order to monitor COD reduction and removal efficiency. The results indicate that
complete reaction was achieved in the first 20 min of reaction as discussed in the
effect of reaction time on Fenton reaction. Consequently, it is impossible to conduct
the kinetic model with two data points of TCOD concentration reduction. Therefore,
the experimental runs at the narrow reaction time of 5, 10, 15 and 20 min were
conducted at the initial pH of 3, [Fe2+] of 0.05 M, and H2O2:Fe2+ molar ratio of 40:1.
The reaction seemed to be fast and completed in about 10 min and the experimental
data on TCOD concentration reduction and its removal efficiencies at different
reaction time are shown in table 4-below. The experimental data then were fitted
according to the integrated forms of the zero, first, second order kinetic models. The
kinetic parameters in kinetic model for Fenton reaction are presented in table 4- . It is
clearly indicated that the first-order kinetic model of Fenton reaction was also fitted to
the experimental data with the good R-square value of 0.9998 as shown in figure 4- .
lower R-square value of the zero-order kinetic model and the second-order kinetic
model were of 0.9240 and 0.9130 respectively. Therefore, this experiment shows that
the rate of Fenton-like reaction is closer to the first-order kinetic model.
Table 4- Change of TCOD concentration with time after Fenton reaction
Time (min)
TCOD (mg/L)
Removal efficiency (%)
0
22157
0.00
5
7296
67.22
10
2268
89.81
15
1936
91.30
20
1935
91.30
80
(a)
(b)
(c)
Table 4- Estimation of the kinetic parameters for Fenton reaction by (a) zero-order,
(b) first order, and (c) second order kinetic models
81
According to the reaction 4-30, the kinetic removal of COD was actually
written in second order reaction in respects to of OH• and COD concentrations
according to the study of Samet et al.,(2011) as provided below:
−
dCOD
= k[OH•]COD
dt
However, Lucas & Peres (2007) and Samet et al.(2011) assume that the
hydrogen peroxide (H2O2) during the reaction is far excess and •OH concentration is
constant during the reaction. Therefore, the kinetic removal of COD during Fentonlike reactions can be simplified to a pseudo-first order reaction as follow:
−
dCOD
= 𝑘𝑎𝑝𝑝 COD  ( pseudo-first order reaction)
dt
where kapp is the pseudo-first order apparent rate constant (kapp = k[OH•])
In this study, •OH concentration is assumed to be constant during the
Fenton-like reaction because excess hydrogen peroxide (H2O2) during the reaction is
applied according to the previous studies (Lucas & Peres, 2007; Wang, 2008; Samet
et al., 2011; Wang et al., 2011). In addition, the first-order kinetic model was
previously carried out without taking consideration of OH• concentration. Therefore,
the obtained first-order kinetic model was actually considered to be the pseudo-first
order model.
In summary, the kinetics of organic matter removal by OH• monitored term
of TCOD reduction for Fenton-like and Fenton reactions can be expressed by pseudofirst order kinetics. At the optimum condition, the apparent kinetic reaction rate
constants of Fenton-like reaction and Fenton reaction were 0.0218 min-1 and 0.2284
min-1, respectively. The kinetic contestant of Fenton reaction is comparatively 10
times higher than Fenton-like reaction. This apparently proves that Fenton reaction
attains faster reaction rate than Fenton-like reaction. The results in this study are in
agreement and comparable with previous studies (Wang et al., 2008; Lucus & Peres,
2009; Wang et al., 2011; Samet et al., 2011; Wu and Wang, 2012) who found Fenton
82
and Fenton-like reactions followed the pseudo-first order reaction. Specifically, Wang
et al., (2008) obtained kinetic constants (k) of 0.0556 min-1 and 0.0311 min-1 for
Fenton reaction and Fenton-like reaction, accordingly for decolourisation of
wastewater. Similarly, Lucas and Peres (2009) conducted kinetic study on COD
removal (initial COD= 60.5 g/L) from olive mill wastewater by Fenton reagent and
found pseudo-first order constant (kapp) of 0.101 min-1 at initial pH of 3.5, molar ratio
H2O2: Fe2+ of 15:1 and weight ratio COD/H2O2 of 1.75. In addition, Wang et al.
(2011) and Samet et al. (2011) similarly obtained kapp of 0.05704 min-1 and 0.07 min1
, respectively.
Comparison between Fenton-like and Fenton reactions
83
CHAPTER 5
CONCLUSION
84
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APPENDICES
Appendix A: Experimental data
Appendix B: Chemical analysis procedures
1.
CHEMICAL OXYGEN DEMAND (COD) (5220 B. Close Reflux Method)
1.1. Principle
Chemical oxygen demand (COD) is used as a measure of oxygen requirement
of a sample that is susceptible oxidation by strong chemical oxidant. The
dichromate reflux method is preferred over procedures using other oxidants
(eg potassium dichromate) because of its superior oxidizing ability,
applicability to a wide variety of samples and ease of manipulation. Most
types of organic matter are oxidized by a boiling mixture of chromic and
sulfuric acids. A sample is refluxed in strongly acid solution with a known
excess of potassium dichromate (K2Cr2O7). After digestion, the remaining
unreduced K2Cr2O7 is titrated with ferrous ammonium sulfate to determine the
amount of K2Cr2O7 consumed and the oxidizable matter is calculated in terms
of oxygen equivalent.
1.2. Apparatus
a. Test tube or digestion vessels
b. Heating oven (150 °C)
c. Micropipettes
d. Beakers (250 mL)
e. Volumetric pipettes (10 mL)
f. Burette and burette stand
g. Test tube rack
1.3. Reagents
93
a. Sulfuric acid reagent (H2SO4) for COD: Add 22g of reagent grade silver
sulphate to a 4-kg bottle of concentrated sulphuric acid (H2SO4) and mix
until the silver sulphate goes into solution.
b. Potassium dichromate (K2Cr2O7) 0.25N, Dissolve 12.259 g K2Cr2O7,
primary standard grade, previously dried at 150°C for 2 h, in distilled
water and dilute to 1000 mL.
c. Ferroin indicator solution: Dissolve 1.485 g 1,10-phenanthroline
monohydrate and 695 mg FeSO4·7H2O in distilled water and dilute to 100
mL.
d. Standard ferrous ammonium sulfate (FAS) titrant, 0.25M: Dissolve 98 g
Fe(NH4 )2(SO4)2·6H2O in distilled water. Add 20 mL conc H2SO4, cool,
and dilute to 1000 mL. Standardize this solution daily against standard
K2Cr2O7 solution as follows:
Dilute 25.00 mL standard K2Cr2O7 to about 100 mL. Add 30 mL conc
H2SO4 and cool. Titrate with FAS titrant using 0.10 to 0.15 mL (2 to 3
drops) ferroin indicator.
Morality of FAS solution=
volume K2 Cr2 O7 solution titrated, mL
×0.1
Volume FAS used in Titration, mL
1.4. Procedure
a. Prepare COD test tubes for sample, hot blank and cold blank
b. Pipette 1 mL of sample and 5 mL of water into sample tube
c. Add 3 mL potassium dichromate
d. Add 7 ml of sulfuric acid reagent
e. Heat in the oven at 150 °C for 2 hours then cool down the test tube
f. Add 2-3 drops of ferroin indicator
g. Start to titrate the FAS solution until sample color turn to mild pink and
record its volume used
(Hot and cold blank tubes are done the same as sample tube processes)
1.5. Calculation
94
COD as mg O2 /L =
(A-B)×M×8000
mL of sample
where:
A = mL FAS used for blank,
B = mL FAS used for sample,
M = molarity of FAS, and
8000 = milliequivalent weight of oxygen X 1000 mL/L.
2. Determination of Iron concentration (3500-Fe B. Phenanthroline Method)
2.1. Principle
Any solution which is colored or can be made colored by adding a complexing
agent can be analyzed using a visible spectrophotometer. Solutions containing
iron ions are colorless, but upon addition of, the iron (II) ions in the sample
react immediately to produce a complex ion, which is orange-red in color. This
follows the Beer-Lambert Law of spectroscopy. From data obtained from a
series of iron (II) standards, it is possible to be able to determine the amount of
iron in an unknown sample.
2.2. Apparatus
a. Spectrophotometer Instrument with its kit; (Varian)
b. Pipette (10 mL)and Micropipette
c. Volumetric Flask (100, 500 mL)
d. Acid-washed glassware
e. Erlenmeyer flasks
f. Hot plate
g. Glass beats
2.3. Reagents
a. Hydrochloric acid (HCl) conc, containing less than 0.5 ppm
b. Hydroxylamine solution: Dissolve NH2OH-HCl in 100 mL water
95
c. Ammonium acetate buffer solution: Dissolve 250 g of NH4C2H3O2 in 150
mL water. Add 700 mL glacial acetic acid. Sodium acetate solution.
d. Sodium acetate solution: Dissolve 200 g NaC2H3O23H2O in 800 mL water
e. Phenanthroline solution: Dissolve 100 mg 1,10-pehnanthroline
monohydrate, C12H8N23H2O in 100 mL water by stirring and heating to
80 °C. Heating is unnecessary if 2 drops of HCL are added.
f. Potassium permanganate (KMnO4), 0.02M: Dissolve 0.316 g KMnO4 in
reagent water and dilute to 100 mL.
g. Stock iron solution: slowly add 20 mL conc H2SO4 to 50 mL water and
dissolve 1.404 g Fe(NH4)2(SO4)23H2O. Then slowly add
potassiumpermanganate solution until fain pink color persists. Add few
milliliters dropwise. Dilute to 1000 mL with water.
h. Standard iron solution: pipette 50 mL stock solution into a 1000 mL
volumetric flash and dilute to mark with water; 1 mL =10 µg Fe.
2.4. Sample preparation before analysis (Sample containing organic
interferences)
2.4.1. Digestion with (HNO3) to digest organic interferences
a. Transfer 50 mL of sample into a 125 mL conical flash
b. Add 5 mL of HNO3conc.
c. Heating on hotplate to evaporate sample to 15-20 mL
d. Heat to remove all HNO3 before continuing treatment (HNO3
removed if solution is clear or no browndish fume is evident
e. Cool down and transfer sample to a 100-mL volumetric flaskdilute
to mark and mix thoroughly. Take 1 portions of this solution for
required metal determinations.
2.5. Determination Procedure
5.2.1. Total iron ( for digested and extracted sample)
a. Take extracted sample (50 mL in a 100-mL volumetric flask) and
add 1 mL of hydroxylamine
96
b. Add 10 ml phenanthroline
c. Add 10 ml Sodium Acetate buffer solution (NaC2H2O2)
d. Dilute to 100 mL then mix and stand for 10 min for color
development
e. Measure absorbance in spectrophotometer and read it concentration
in calibration curve (figure 1)
f. Calibration curve of ferrous iron is formulated by pipetting 50 mL of
standard solution (1mL = 10 µg Fe) in a 100 mL volumetric flask
then follow step a-c (Series dilutions are used to get ferrous iron
concentration of 31.25, 62.5, 125, 250, and 500 µg in final 100volume)
Figure 1: Calibration curve for total iron
5.2.2. Dissolved iron
a. Immediately collect filter sample through a 0.45 µm filter paper into
a vacuum flask containing 1 mL conc HCl/100 mL sample. Analyze
97
filtrate for total dissolved iron (according to step 5.2.1) or ferrous
iron( according to step 5.2.4)
Suspended Iron = Total Iron- Dissolved Iron
5.2.3. Ferrous iron:
a. Acidify the sample by pipetting 50 mL into 125-mL Erlenmeyer
flask (if Fe >200 µg, the sample needs to be diluted) and add 1 mL
conc HCl
b. Add 20 mL phenanthroline solution and 10 mL NH4C2H3O then
Dilute to 100 mL
c. Measure color intensity within 5 to 10 min in spectrophotometer at
510 nm and concentration of ferrous from calibration curve as
provided in figure 1
d. Ferric iron can be determined by
Ferric Iron = Total Iron-Ferrous Iron
5.3. Calculation
mg Fe/L=
µg Fe (in 100 mL final volume)
50
×
mL sample
ml portion
Appendix C: Fenton’s reagent preparation
1. Preparation of 300 mL of Fe2+ solution (0.1M) from FeSO4.7H2O
a. Calculate the amount of Fe2+ from FeSO4.7H2O by
250 mL× 0.1
mole Fe2+ 1 mol FeSO4 .7 H2 O
278.02 g
×
×
×
L
1 mol FeSO4 .7 H2 O
1 mol Fe2+
1L
100
(% Assay)= 69.8542 g
×
1000mL 99.5
b. Weight 69.8542 g of FeSO4.7H2O using analytical balance
98
c. Dissolve 69.8542 g of FeSO4.7H2O and transfer to 250 mL volumetric flash and
dilute to the mark
d. Adjust pH to 3-4 in order to avoid the precipitation of iron then keep it in
refrigerator for the next experiment
2. Calculation of H2O2 molar concentration
2.5 L of H2O2 with assay 35 % (v/v) is used in the experiment. Density of 35% H2O2
= 1.14 g/mL and MW of H2O2 = 34.01 g/mol
Density × % of H2 O2 ×
1.14
1 mol
1000 mL
×
MW of H2 O2
1L
g
35
1 mol
1000 mL
×
×
×
=11.73 M
mL 100 34.01 g
1L
3. Preparation of 10 N NaOH Solution
a. Calculate the amount of NaOH required
10 eq mol
40 g
×
×
= 400 g/L
L
1 eq
mol
b. If solution volume is required 500 mL, mass of NaOH required is 200 g
c. Dissolve 200 g with distilled water then dilute to 500 mL
4. Preparation of 5 N H2SO4 Solution
a. Calculate the concentration of H2SO4 ( density= 1.84 g/cm3 or 1.84 g/mL, purity=
98%, MW=98.08 g/mol)
Density× % of H2 SO4 ×
1 mol
1000 mL
×
MW of H2 SO4
1L
1.84 g 98
1 mol
1000 mL
×
×
×
= 18.38 M
mL
100 98.08 g
L
Normality (N) of H2SO4 is given by
99
18.38 mol 2 eq
×
= 36.76 M
L
1 mol
b. Calculate the volume of H2SO4 for 500 mL of H2SO45N
According to dilution rule: N1V1= N2V2
36.76N  V1 = 5N 0.5L
The required volume of H2SO4 = 0.068 L or 68 mL
c. Prepare 250 mL of distilled water in a 500 mL beaker
d. Pipette 68 mL of con. H2SO4 and add slowly into 500-mL beaker then dilute to
the mark
e. Cool down and ready to use for pH adjustment
Appendix D: Conceptual experimental pictures