Photocatalytic Degradation of Light Green Dye Using

Republic of Iraq Ministry of Higher Education and Scientific Research University of Kufa - Faculty of Science Chemistry Department

Photocatalytic Degradation of Light Green Dye Using TiO2 and Nano TiO2 as Catalysts A thesis Submitted to the Council of the Faculty of Science University of Kufa as a Partial Fulfillment of the Requirements for M.Sc. in Chemistry

By Mohammed Taher Eassa B.Sc. 2012 (University of Karbala)

Supervised by Assist.prof.Dr. Amer Mousa Juda

2016 A

Assist.prof.Dr. Luma Majeed Ahmed

1438 H

‫اﻟﻘﺮآن اﻟﻜﺮﯾﻢ‪ :‬ﺳﻮرة اﻟﻔﺎﺗﺤﺔ‬

Dedication

To my dear parents, who are the reason behind my existence and success

To my beloved wife, who supported me along the way

To my dear son …

To my dear friends …

Mohammed 2016

Acknowledgement

Initially, I want to thank the almighty Allah for giving me the strength and guidance throughout my entire life and during this work in particular. I wish also to express my deepest gratitude to my supervisors Asst. Prof. Dr. Amer Mousa Jodah and Asst. Prof. Dr. Luma Majeed Ahmed for their continuous support and invaluable suggestions and great contributions since the very beginning of this work. Also, I want to thank all faculty members of Department of Chemistry in the Faculty of Science at The University of Kufa, for their irreplaceable support. Also, I want to thank all faculty members of Department of Chemistry at the Faculty of Science at The University of Karbala, for their worthless support of the work. Finally, I want to thank my family and friend for their continual support throughout this journey. They were my source of encouragement along the way.

Mohammed

Abstract This work consists of three parts: Part one: Includes the preparation of titanium dioxide using sol-gel method by the reaction between titanium tetrachloride and ethanol with different mixing ratios (1:4, 3:10, 1:10), and under two different calcination temperatures (600, 800 °C). Part two: Includes the characterization of commercial bulk and commercial nano and prepared TiO2 by X-ray diffraction (XRD), Atomic force microscopy (AFM) and Fluorescence spectroscopy techniques. The obtained XRD patterns substantiate that no rutile or brookite phase present in the samples calcinated at 600 °C, and that rutile phase appeared in the samples calcinated at 800°C. Crystallite size and average crystallite size of the commercial bulk, commercial nano and prepared TiO2 were calculated employing Scherrer and modified Scherrer equations. AFM images showed that particles of the prepared titanium dioxide have spherical shape. Fluorescence spectroscopy showed that the band gap energy of the studied catalysts is in the range (3.51 – 3.56 eV) which is slightly larger than reported values. Part three: This part investigated the effect of calcination temperature on the photocatalytic activity of the prepared samples TiO2, and it was found that the optimum catalyst among the samples was that with mixing ratio (3:10) calcined under 600 °C. This part investigated also various factors affecting the photocatalytic degradation of Light Green SF Yellowish dye in the presence of the commercial bulk, commercial nano and prepared TiO2. The studied factors were the initial

I

Abstract

concentration of LGSF, amount of catalyst, initial pH of the LGSF solution and the temperature of the dye solution. The effect of the initial concentration of the dye solution was studied for different concentrations in the range (20-50 ppm), and the results showed that the reaction is a pseudo first-order reaction. The optimal dosage of the catalyst was found to be (0.3 g/100 mL) for the commercial bulk TiO2, and (0.2 g/100 mL) for both commercial and prepared nano TiO2. Investigation of the effect of initial pH of the dye revealed that the optimum pH of the solution is (7.3) for each type of the catalysts used in this work. Effect of temperature was investigated also employing Arrhenius equation, and the results showed that the increase in temperature is accompanied with an increase in the rate of the reaction in the range (278.15 – 293.15 K), indicating that the photocatalytic degradation of Light Green dye is an endothermic reaction. The activation energy of the reaction was calculated for each type of the catalyst, and it was found to be (48.676) kJ.mol-1 for commercial bulk TiO2, and (37.032) kJ.mol-1 for commercial nano TiO2, (45.780) kJ.mol-1 for the prepared TiO2. Entropy and enthalpy values were calculated employing Eyring-Polanyi equation. Entropy values were found to be (– 0.107) kJ.mol-1 for commercial bulk TiO2, and (– 0.152) kJ.mol-1 for commercial nano TiO2, (– 0.120) kJ.mol-1 for the prepared TiO2, these results showed that the randomness was decreased. Enthalpy values were found to be (46.292) kJ.mol-1 for commercial bulk TiO2, and (34.667) kJ.mol-1 for commercial nano TiO2, (43.416) kJ.mol-1 for the prepared TiO2. These results show clearly that the photocatalytic degradation of LGSF dye is an endothermic process. Gibb’s free energy was calculated for each type of the photocatalysts used in this study, and the values were found to be (78.847) kJ.mol-1 for commercial bulk TiO2, and (80.664) kJ.mol-1 for commercial nano TiO2, (79.707) kJ.mol-1 for the prepared TiO2. These results show that the reaction is non-spontaneous.

II

Table of Contents Content

Page

Abstract

I

Table of Contents

III

List of Tables

VIII

List of Figures

IX

List of Abbreviations and symbols

XVI

Chapter One: Introduction 1.1

General Introduction

1

1.2

Nanochemistry

2

1.2.1

Classification of nanomaterials

2

1.2.2

The effect of size

3

1.2.3

Physical properties of nanomaterials

5

1.2.3.1 Geometric structure

6

1.2.3.2 Optical properties

6

1.2.3.3 Thermal properties

6

1.2.3.4 Electric properties

6

1.3

Methods of nanoparticles synthesis

6

1.4

Advanced oxidation processes

7

III

Table of Contents

1.4.1

Advantages and disadvantages of AOPs

12

1.5

Photochemistry

12

1.5.1

Photocatalysis

14

1.5.1.1 Photocatalytic degradation

15

1.5.1.2 Photocatalytic degradation of dyes

15

1.5.2

Quantum yield

17

1.6

Semiconductors

17

1.6.1

Classification of semiconductors

20

1.6.1.1 Intrinsic semiconductors

21

1.6.1.2 Extrinsic semiconductors

21

1.7

Titanium Dioxide

22

1.7.1

Photolysis effect

24

1.7.2

TiO2 nanoparticles

26

1.7.3

Sol-gel method

27

1.8

Adsorption

29

1.8.1

Adsorption on semiconductors surface

29

1.8.2

Adsorption of oxygen

31

1.8.3

Adsorption of dyes

32

1.9

Dyes

33

1.9.1

Triarylmethane dyes

34

1.9.2

Light Green dye

35

1.10

Aim of the study

36

IV

Table of Contents

Chapter Two: Experimental part 2.1

Chemicals

37

2.2

Instruments

38

2.3

Preparation of TiO2 nanoparticles by sol-gel method

39

2.4

X-ray diffraction patterns

41

2.5

Atomic force microscopy (AFM)

42

2.6

Band gap energy measurement

42

2.7

Light intensity measurement

43

2.8

Measurement of maximum absorption wavelength of the dye

45

2.9

Calibration curve of the dye

46

2.10

Measurement of the optimal conditions

47

2.10.1

Optimal concentration

48

2.10.2

Optimal catalyst dose

48

2.10.3

Optimal pH

49

2.10.4

Optimal temperature

49

2.11

Photocatalytic kinetic analysis

49

2.12

Photocatalytic dcolorization efficiency

50

2.13

Activation energy

51

2.14

Thermodynamic parameters

51

Chapter Three: Results 3.1

Catalysts characterization

52

3.1.1

X-Ray diffraction patterns (XRD)

52

3.1.2

Atomic force microscopy (AFM)

57

V

Table of Contents

3.1.3

Band gap energy measurements

66

3.2

Photocatalytic degradation of LGSF dye

68

3.2.1

Preliminary experiments

68

3.2.1.1 Dark reaction

68

3.2.1.2 Photolysis

69

3.2.2

Effect of different parameters on the photocatalytic degradation of LGSF dye for commercial TiO2

70

3.2.2.1 Effect of initial concentration of the dye

70

3.2.2.2 Effect of the mass of commercial TiO2

71

3.2.2.3 Effect of the initial pH of the solution

73

3.2.2.4 Effect of the temperature

74

3.2.3

Effect of different parameters on the photocatalytic degradation of LGSF dye for commercial nano TiO2

77

3.2.3.1 Effect of initial concentration of the dye

77

3.2.3.2 Effect of the mass of commercial nano TiO2

78

3.2.3.3 Effect of the initial pH of the solution

80

3.2.3.4 Effect of the temperature

81

3.3

Effect of calcination temperature on prepared TiO2

84

3.4

Effect of different parameters on the photocatalytic degradation of LGSF dye with prepared TiO2 and calcination at 600 °C

86

3.4.1

Effect of initial dye concentration

86

3.4.2

Effect of the mass of prepared TiO2

87

3.4.3

Effect of initial pH of the solution

89

3.4.4

Effect of temperature

90

VI

Table of Contents

Chapter Four: Discussion 4.1

Characterization of the catalysts

93

4.1.1

X-Ray diffraction spectroscopy (XRD)

93

4.1.2

Atomic force microscopy

93

4.1.3

Fluorescence spectroscopy

93

4.2

Preliminary Experiments

94

4.3

Effect of different parameters on photocatalytic degradation of LGSF dye using commercial, commercial nano and 95 prepared TiO2

4.3.1

Effect of initial dye concentration

95

4.3.2

Effect of the mass of the catalyst

96

4.3.3

Effect of pH on dye removal

98

4.3.4

Effect of temperature on dye removal

100

4.4

Suggested mechanism

102

Chapter Five: Conclusions, Recommendations and References 5.1

Conclusions

104

5.2

Recommendations

105

References

106

Appendix A

117

Appendix B

125

Appendix C

132

VII

List of Tables Table 1-1

Title of Table

Page

Heats of formation of AOPs products and some organic

8

molecules. 1-2

Classification of Advanced Oxidation Processes.

9

1-3

Comparison of rate constant for ozone and hydroxyl radicals for

10

a range of compounds. 1-4

Standard reduction potentials of some oxidants in volts (V)

11

compared with the normal hydrogen electrode (NHE, E0 = 0 V). 1-5

Differences between photochemical and thermochemical reactions.

13

1-6

Comparison between p-type and n-type semi-conductors.

22

1-7

Crystal structure data for TiO2.

25

1-8

Classification of dyes according to chemical structure.

33

2-1

Chemicals and their commercial sources.

37

2-2

Instruments used in this work.

38

2-3

Absorbance at different concentrations of Fe(II).

45

2-4

Absorbance at different concentrations of LGSF.

46

3-1

Crystallite size and average crystallite size of commercial bulk,

57

commercial nano and prepared TiO2 3-2

Particles size of commercial bulk, commercial nano and

66

prepared TiO2 samples. 3-3

Band gap energy measurements results.

VIII

68

List of figures Figure

Title of figure

Page

1-1

Classification of nanomaterials depending on dimension-ality.

3

1-2

The relationship between surface area (SA) and volume (Vol)

4

with respect to the size of the material. 1-3

Proportion of surface atoms for a spherical particles

5

comprising NT atoms with Ns at the surface. 1-4

The three methods of nanoparticles synthesis.

7

1-5

AOP technologies for water and air purification.

9

1-6

Schematic representation of the mechanisms of generation of

17

oxidative species in photocatalytic degradation technique. 1-7

Electronic energy levels in a bulk semiconductor.

18

1-8

Position and width of the energy band of TiO2 and several

19

other illuminated semiconductors with respect to the electrochemical scale (NHE: normal hydrogen electrode). 1-9

A sketch of the energy bands for a semiconductor with direct

20

and indirect bandgap. 1-10

The classification of semiconductors.

21

1-11

Unit cells of TiO2 (a)rutile, anatase(b) and brookite(c).

23

1-12

Synthesis methods of TiO2 nanostructures.

26

1-13

Schematic diagram for Sol-gel method.

27

1-14

Adsorption of oxygen on the surface of TiO2.

31

1-15

Classification of dyes according to application method.

33

1-16

The chemical structure of Fuchsine.

34

1-17

Chemical structure of LGSF.

35

IX

List of Figures

2-1

Synthesis setup.

39

2-2

Detailed schematic representation for experimental procedure.

40

2-3

Schematic diagram for experimental procedure

41

2-4

The experiment of light intensity measurement.

43

2-5

Calibration curve for Fe(II) as a complex with 1,10-

45

phenanthroline. 2-6

UV-Visibile spectrum of LGSF dye squeous solution.

46

2-7

Calibration curve at different concentrations of LGSF Dye

47

2-8

The photocatalytic reactor.

48

3-1

X-ray diffraction spectrum for the commercial TiO2.

52

3-2

X-ray diffraction spectrum for the commercial nano TiO2.

52

3-3

X-ray diffraction spectrum prepared TiO2 nanoparticles with

53

600 °C calcination temperature. [V/V (TiCl4:EtOH) = 1:10]. 3-4


53



54



54



55



55


Modified Scherrer equations of commercial bulk, commercial

56

nano and prepared nano TiO2. 3-10

AFM images of the commercial TiO2.

58

3-11

AFM images of the commercial nano TiO2.

59

X

List of Figures

3-12

AFM images of the prepared TiO2 nanoparticles with 600°C

60

calcination temperature. [V/V (TiCl4:EtOH) = 1:10]. 3-13

AFM images of the prepared TiO2 with 800°C calcination

61

temperature. [V/V (TiCl4:EtOH) = 1:10]. 3-14


62



63



64



65


Fluorescence spectrum of commercial bulk TiO2.

66

3-19

Fluorescence spectrum of commercial nano TiO2.

67

3-20

Fluorescence spectrum of prepared TiO2 nanoparticles with

67

600°C calcination temperature [V/V (TiCl4:EtOH) = 3:10]. 3-21

The variation of ln(C0/Ct) with time in the absence of radiation.

69

3-22

The variation of ln(C0/Ct) with time in the absence of the

69

catalyst. 3-23

The change of ln(C0/Ct) with irradiation time for different

70

initial concentrations of LGSF dye solution using commercial TiO2. 3-24

The relationship between PDE and irradiation time for different

71

initial concentrations of LGSF dye solution using commercial TiO2. 3-25

The change of ln(C0/Ct) with irradiation time for different masses of commercial TiO2.

XI

72

List of Figures

3-26

The relationship between PDE and irradiation time for

72

different masses of commercial TiO2. 3-27


73

initial pH value of the dye solution in the presence of commercial TiO2. 3-28


74

different initial pH of LGSF dye solution using commercial TiO2. 3-29


75

temperatures in the presence of commercial TiO2. 3-30


75

different temperatures using commercial TiO2. 3-31

Arrhenius relationship with commercial TiO2.

76

3-32

Eyring plot of ln(k/T) vs. 1/T.

76

3-33


77

initial concentrations of LGSF dye solution using nano commercial TiO2. 3-34


78

different initial concentrations of LGSF dye solution using nano commercial TiO2. 3-35

Variation of ln(C0/Ct) with irradiation time for different

79

masses of commercial nano TiO2. 3-36


79

different masses of commercial nano TiO2. 3-37

The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of commercial nano TiO2.

XII

80

List of Figures

3-38


81

different initial pH of LGSF dye solution using commercial nano TiO2. 3-39


82

temperatures in the presence of commercial nano TiO2. 3-40


82

different temperatures using commercial nano TiO2. 3-41

Arrhenius relationship with commercial TiO2.

83

3-42

Eyring plot of ln(k/T) vs. 1/T.

83

3-43


84

mixing ratios of the prepared TiO2 calcinated under 600 °C. 3-44


84

mixing ratios of the prepared TiO2 calcinated under 800 °C. 3-45


85

different mixing ratios of the prepared TiO2 calcinated under 600 °C. 3-46


85

different mixing ratios of the prepared TiO2 calcinated under 800 °C. 3-47

The variation of the apparent rate constant with the variation

85

of the mixing ratio at 600 and 800° calcination temperatures 3-48

The change of ln(C0/Ct) with irradiation time for different initial concentrations of LGSF dye solution using prepared TiO2.

XIII

86

List of Figures

3-49


87

different initial concentrations of LGSF dye using prepared TiO2. 3-50

The variation of ln(C0/Ct) with irradiation time for different

88

masses of the prepared TiO2. 3-51


88

different masses of prepared TiO2. 3-52

The variation of ln (C0/Ct) with irradiation time for different

89

initial pHs of the dye solution. 3-53


90

different initial pHs of the dye solution. 3-54

The change of ln(C0/Ct) with irradiation time at different

91

temperatures using prepared TiO2 as a catalyst. 3-55

The relationship between PDE and irradiation time at different

91

temperatures using prepared TiO2 as a catalyst 3-56

Arrhenius relationship with prepared TiO2.

92

3-57

Eyring plot with prepared TiO2.

92

4-1

Preliminary experiments with commercial bulk, commercial

94

nano and prepared TiO2. 4-2

The relationship between the apparent rate constant and the

96

initial concentration of LGSF dye. 4-3

The relationship between the apparent rate constant and the

97

dosage of the catalyst. 4-4

The relationship between the apparent rate constant and pH of dye solution.

XIV

99

List of Figures

4-5

The relationship between the apparent rate constant and pH of

100

dye solution. 4-6

The relationship between lnk and 1/T.

102

4-7

The relationship between ln(k/T) and 1/T.

102

4-8

Suggested mechanism for the degradation of LGSF dye.

103

XV

List of Abbreviations and Symbols Abbreviation

Description

AFM

Atomic Force Microscopy

AOPs

Advanced Oxidation Processes

CB

Conduction Band

C0

Initial concentration

Ct

Concentration of the substrate at time (t) of irradiation

e–

Negative electron

Ea

Activation energy

Eg

Energy gap

eV

Electron Volt

h+

Positive hole

I0

Light intensity

L

Average crystallite size

Ĺ

Crystallite size

LGSF

Light Green SF Yellowish

PDE

Photocatalytic decolorization efficiency

pHpzc

Point of zero charge

SPM

Scanning Probe Microscopy

rpm

Rotation per minute

UV

Ultraviolet

UV(A)

Ultraviolet light in the range from 315 to 380 nm

VB

Valence band

XRD

X-Ray diffraction

μ

Refractive index

XVI

Chapter One

Introduction

Chapter One: Introduction

1.1. General Introduction Each year a significant amount of wastewater is produced from industries like textile, paper, dyestuffs and plastics. Among these contaminants dyes are considered the primary pollutant in wastewater[1]. From an ecological point of view the presence of these dyes, even in very small concentrations (below 1 ppm), is undesirable since it blocks the sunlight which is essential for photosynthesis of aquatic plants. In the course of the manufacturing process 2% of the dyes is discharged directly in aqueous effluents, and 10% of the dye loss occurs during the textile coloring processes[2]. Therefore, finding an effective method to remove these dyes from wastewater is a necessity, and various physical and chemical processes, like adsorption on activated carbon and chemical precipitation are currently used. However, these methods are nondestructive methods and the contaminants are transferred from one phase to another, which means that further treatment may be required to get rid of these pollutants[3]. Photocatalysis, which is an AOP process that uses semiconductors like TiO2 and ZnO to degrade organic pollutants, is used in the recent years as an efficient alternative for the treatment of wastewater polluted with toxic organic compounds such as dyes[4]. Nanotechnology, which promises to provide a wide range of improved technologies and novel uses for various applications, has found its way to photocatalysis through the production of semiconductor nanoparticles leading to a great enhancement and development in this technique.

1


1.2. Nanochemistry The term nanochemistry was invented originally by Ozin in 1992[5], and it is composed of the two words “nano” and “chemistry”. “Nano” is a prefix used in the metric system to denote 10-9[6], and it stems from the Greek word “nanos” which means dwarf[7]. Nanomaterials refer to class of materials with at least one of their dimensions in the nanometric range (< 100 nm), and they can be metals, polymers, ceramics or composites[8], however, sometimes this science is defined in a more restricted manner as the study of new effects that appear merely in materials that exist on the nanoscale[9].

1.2.1. Classification of nanomaterials Presently, the most acceptable way of classifying nanomaterials is to identify them according to their dimensions. According to Seigel Nanomaterials are classified into four types based on the number of dimensions, which are not restricted to the nanoscale range (< 100 nm)[10]: i.

Zero dimensional (0D): Materials wherein all the dimensions (x,y,z) are measured within the nanoscale, i.e. no dimensions are larger than 100 nm[11]. Many types of 0D nanomaterials have been synthesized by several research groups, such as uniform particle arrays (quantum dots), heterogeneous particles arrays, core–shell quantum dots and hollow spheres[12].

ii.

One dimensional (1D): This class of nanostructures has nanoscale sizes along two dimensions and has a rod-like or wire-like appearance. 1D nanostructures have a large impact in nanoelectronics, nanodevices and systems, nanocomposite materials and alternative energy resources[13].

2


iii.

iv.

Two dimensional (2D): 2D nanostructures have two dimensions outside the nanometric size range. This class comprises different types of nanomaterials such as nanoprisms, nanoplates, nanosheets and nanodisks[14]. Three dimensional (3D): These materials are known as bulk nanomaterials and have no dimension that is confined to the nanoscale range[15]. Despite, their nanoscale dimensions bulk nanomaterials possess nanocrystalline structure or involve the presence of features at the nanoscale range (can contain dispersions of nanoparticles, bundles of nanowires, and nanotubes).

Figure 1-1: Classification of nanomaterials depending on dimensionality[10].

1.2.2. The effect of size One of the most fundamental differences between nanomaterials and largerscale materials is that nanoscale materials have an extraordinary ratio of surface area to volume. For a spherical material of radius (r) then: Surface area of thesphere 4πr 2 3 = = Volume of thesphere 4 3 r  3  πr  

3

1-1


The above ratio increases dramatically as the radius of the spherical material decreases. Thus, the surface area to volume ratio is larger in nanomaterials than bulk materials[16]. Figure 1-2 depicts the relationship between surface area and volume with respect to the size of the material.

Figure 1-2: The relationship between surface area (SA) and volume (Vol) with respect to the size of the material. Though the properties of traditional large-scale materials are often determined entirely by the properties of their bulk, due to the relatively small contribution of a small surface area, for nanomaterials this surface-to-volume ratio is inverted, and as a result, the larger surface area of nanomaterials (compared to their volume) plays a larger role in dictating these materials’ important properties. This inverted ratio and its effects on nanomaterials properties is a key feature of nanoscience and nano-technology[15]. When the size of the object is reduced to the nanometric range, i.e. < 10 nm, the proportion of surface atoms to the total number of atoms is no longer

4


negligible. For instance, at 5 nm (around 8,000 atoms) the proportion is about 20%, whilst at 2 nm (around 500 atoms) it shifts to 50%[17].

Figure 1-3: Proportion of surface atoms for a spherical particles comprising NT atoms with Ns at the surface[17]. This proportion can be estimated for the transition metals by the relation

Ns 1 ≈ N T 2r

1-2

Where: r is the radius of particle in nm, NT is the total number of atoms and Ns is the number of atoms at the surface. This empirical law gives a proportion of surface atoms of 100% for a size of 1 nm, and it is obvious that equation 1-2 is not valid for smaller dimensions.

1.2.3. Physical properties of nanomaterials The interaction between atoms in nanostructures is different from that in bulk materials, and this dissimilarity leads to different properties compared to what they exhibit on macroscales[16]. 5


1.2.3.1. Geometric structure Large nanoparticles have different lattice parameters than that of bulk materials, and the surface to volume ratio increases, whereas the distance between atoms decreases due to reduction of the size of nanoparticles[16]. 1.2.3.2. Optical properties Nanoparticles with different sizes scatter the incident light on it at different wavelengths thus they appear with diverse colors. Gold nanoparticles for instance can appear with orange, purple, red or green color depending on the nanoparticle size[18]. 1.2.3.3. Thermal properties Going from bulk to nanoparticles affects the thermal properties of some materials. For example, Stable aluminum becomes combustible in nanophase, and solid gold transforms into liquid at room temperature as it goes from bulk to nanomaterial[19]. 1.2.3.4. Electrical properties The electric properties of the materials also change when going from bulk to nanomaterials. For instance, bulk silicon is an insulator, but it converts into a conductor in nanophase[16].

1.3. Methods of nanoparticles synthesis The methods of nanoparticles synthesis are divided into three types, sketched in Figure 1-4[15, 20]: 1- Top-down: involves breakdown of the structure of a bulk material, reducing the crystal size to submicron or nano-dimensions. 2- Intermediate: starts with standard microscale particles, reducing their structure to the nanoscale by milling techniques.

6


3- Bottom-up: builds up the solid from the atomic scale or from nanoclusters in a way that retains the scale of its structural units. This method of synthesis yields more control in nanoscale synthesis, but usually is more expensive than “top-down” methods.

Figure 1-4: The three methods of nanoparticles synthesis[15].

1.4. Advanced oxidation processes Glaze et al., (1987) defined Advanced Oxidation Processes (AOPs) as the technologies characterized by the generation of highly reactive non-selective radicals such as hydroxyl radicals (●OH) to degrade toxic organic pollutants from wastewaters[4, 21]. Figure 1-5 depicts the advanced oxidation technologies for water and air purification.

7


The ultimate goal of the oxidation of pollutants in water is to ‘mineralize’ i.e. to convert organic pollutants into simple, relatively harmless and inorganic molecules[22]. • Carbon is converted into carbon dioxide. • Hydrogen is converted into water. • Phosphorous is converted into phosphates or phosphoric acids. • Sulfur is converted into sulfates. • Nitrogen is converted into nitrates. • Halogens are converted into halogen acids. The stability of the ultimate products is the driving force for oxidation, since the heat of formation of these products is much less than that of organic molecules. Table 1-1 lists heats of formation of these products. Table 1-1: Heats of formation of AOPs products and some organic molecules. Product

Heat of formation kJ.mol–1

CO2(g)

–393

H2O(l)

–285

PO3– 4(aq)

–1276

SO2– 4(aq)

–908

NO–3(aq)

–207

Cl–(aq)

–167

C2H7N

–74

C6H6

+49

So, for instance, conversion of ethylamine to CO2, H2O and NO–3(aq) is thermodynamically favorable since about 2000 kJ.mol–1 of energy is released[22].

8


Solar Processes Photocatalysis

Photochemical processes

Supe Critical Water Oxidation Electron Beam Irradiation

Catalytic Processes

AOP

Elechtrochemical Processes

γ-Radiolysis X-Ray Irradiation

Non-thermal Plasma Techniques

Sonolysis

Figure 1-5: Schematic description of AOP technologies for water and air purification. Adapted from[23]. AOPs can be classified as homogeneous and heterogeneous[24], or in terms of whether light is used in the process or not. Table 1-2 shows a classification of AOPs according to the two preceding methods. Table 1-2: Classification of Advanced Oxidation Processes[4, 25]. Non-photochemical

Photochemical

Homogenous processes Ozonation in alkaline media (O3/HO–)

Photolysis of water in vacuum ultraviolet (VUV)

Ozonation with hydrogen peroxide (O3/H2O2)

UV/H2O2

Fenton (Fe2+ or Fe3+/H2O2)

UV/O3

Electro-oxidation

UV/O3/H2O2

Electrohydraulic discharge - ultrasound

Photo-Fenton (Fe2+ or Fe3+ /H2O2/UV)

Heterogeneous processes Catalytic wet air oxidation (CWAO)

Heterogeneous photocatalysis: ZnO/UV, SnO2/UV, TiO2/UV, TiO2/H2O2/UV

9


The chemistry of AOPs can be divided into three parts[26]: 1- Formation of strong oxidants (e.g. ●OH… etc). 2- Initial attacks on target molecules by ●OH and their breakdown to biodegradable intermediates. 3- Subsequent attacks by ●OH until ultimate mineralization (i.e. production of water, carbon dioxide and inorganic salts). OH radicals are extraordinarily reactive species; they attack most part of organic molecules with rate constants usually in the order of 106–109 M−1.s−1[27]. OH radicals are able to oxidize a wide range of organic compounds significantly faster than O3 (see Table 1-3)[22]. Table 1-3: Comparison of rate constant for ozone and hydroxyl radicals for a range of compounds[22]. Rate constant (M–1.s–1)

Organic compound

O3

●

OH

Benzene

2

7.8 × 109

Toluene

14

7.8 × 109

Chlorobenzene

0.75

4 × 109

Trichloroethylene

17

4 × 109

Tetrachloroethylene

< 0.1

1.7 × 109

n-Butanol

0.6

4.6 × 109

Hydroxyl radical is a powerful oxidant and a short lived, highly reactive, and non-selective reagent that is easy to produce[23], in addition, it is a highly reactive electrophile; in fact, it is regarded as the second most powerful oxidant after Fluorine as shown in Table 1-4.

10


Table 1-4: Standard reduction potentials of some oxidants in volts (V) compared with the normal hydrogen electrode (NHE, E0 = 0 V)[4, 28]. Oxidant

Standard reduction potential / V

Fluorine (F2)

3.03

Hydroxyl radical (●OH)

2.80

Positively charged hole on titanium dioxide (TiO+2 )

2.35

Ozone (O3)

2.07

Hydrogen peroxide (H2O2)

1.77

Potassium permanganate (KMnO4)

1.67

Chlorine dioxide (ClO2)

1.50

Chlorine (Cl2)

1.36

Bromine (Br2)

1.09

Mechanism for ●OH production depends highly on the AOP technique, for instance: • UV/H2O2[29]: H2 O2 + UV �⎯⎯⎯� 2 ● OH

1-3

O3 + HO– �⎯⎯⎯� HO–2 + O2

1-4

• Ozone based AOP[30]:

O3 + HO–2 �⎯⎯⎯⎯� HO●2 + O–● 3 ● + O–● 3 + H �⎯⎯⎯⎯� HO3

HO●3 �⎯⎯⎯⎯⎯� ●OH+O2

1-5 1-6 1-7

• Semiconductor photocatalysis[31]: hν

TiO2 �⎯⎯� TiO2 (e–CB ) + TiO2 (h+VB )

TiO2 �h+VB �+TiO2 (H2 O) �⎯⎯⎯⎯� TiO2 (HO● )+H+ 11

1-8 1-9


TiO2 (e–CB )+O2 �⎯⎯⎯⎯� TiO2 +O●– 2 ● + O●– 2 +H �⎯⎯⎯⎯� HO2

HO●2 + HO● �⎯⎯⎯� H2 O+O2

Dye+ e–CB �⎯⎯⎯⎯� reduction products Dye+ h+VB �⎯⎯⎯� oxidation products

Dye+ HO● �⎯⎯⎯� degradation product

1-10 1-11 1-12 1-13 1-14 1-15

1.4.1. Advantages and disadvantages of AOPs: The advantages of AOPs can be summarized in the following points: • Fast reaction rates. • Reduce toxicity and mineralize waste materials almost completely. • Does not concentrate waste for further treatment with methods such as membranes. • Does not produce materials that require further treatment • Does not create sludge as with physical chemical process or biological processes (wasted biological sludge). Whereas disadvantages are: • It could be an expensive technology. • Some applications require quenching of excess peroxide.

1.5. Photochemistry Photochemistry is the study of the chemical reactions and physical changes that result from interactions between matter and visible or ultraviolet light[32]. Hence, a photochemical reaction, according to the preceding definition, may be defined as a chemical reaction caused by absorption of light or ultraviolet radiation by reacting species. The incident photons are absorbed by reactant molecules to give excited molecules or free radicals, which undergo further reactions[6].

12


There are significant differences between photochemical and thermochemical reactions, which are given in Table 1-5. Table 1-5: Differences between photochemical and thermochemical reactions[33, 34]. Photochemical Reactions

Thermochemical Reactions

These reactions are initiated by light These reactions are initiated by heat radiation.

energy.

These involve absorption of light

These reactions involve absorption or

Radiations.

evolution of heat.

They cannot occur in dark.

They can occur in dark as well as in light.

Temperature has a very little effect on Temperature has a significant effect on the rate of photochemical reactions.

the rate of a thermochemical reaction.

ΔG for photochemical spontaneous ΔG for a thermochemical reaction is reactions may be positive or negative.

always negative.

Photochemical activation is highly Thermochemical

activation

is

not

selective. The absorbed photon excites selective in nature. a particular atom or group of atoms which become site for the reaction. The excited state of a molecule is The ground state of a molecule is involved in these reactions.

involved in these reactions.

The viability of a photochemical reaction depends on the fulfillment of the following requirements[35]: • The photochemical reaction should be endothermic.

13


• The process should- be cyclic. • There is on side reactions that degrade the photochemical reactants. • The reaction should use the widest possible range of the solar spectrum. • The quantum yield should approach unity. • Under normal conditions, back reaction should be slow allowing the storage of the products, but it should be rapid under specific controlled conditions to release the energy. • The reagents and container material should be inexpensive and non-toxic. • The process should operate under aerobic conditions.

1.5.1. Photocatalysis The word “photocatalysis” is of Greek origin and composes of two parts: “photo” from the Greek word phōs (light)[7] and “catalysis” which was coined in 1836 by Berzelius from the Greek words kata (wholly) and lyein (to loosen) [36]. This term is used to describe a process in which a photochemical reaction is accelerated by the presence of a catalyst that absorbs light and is involved in the chemical transformation of the reactants[37]. Photocatalysis can be divided into two classes[36]: 1- Homogenous photocatalysis: all the substances involved in the reaction including the photocatalyst are in the same phase. The most commonly used homogenous photocatalysts include ozone and photo-Fenton systems (Fe+ and Fe+/H2O2). 2- Heterogeneous photocatalysis: the photocatalyzed reaction occurs at the boundary between two phases, and usually on the surface of a solid photocatalyst. For a unimolecular process in the gas phase, for instance:

14


A(gas) �⎯⎯⎯⎯� products

1-16

The catalyzed mechanism is[36]: A(gas) + surface site

hv

A(adsorbed)

A(adsorbed) �⎯⎯⎯⎯� products

1-17 1-18

A good photocatalyst should have the following properties[38]: 1- Photoactive. 2- Able to absorb visible and/or UV light. 3- Chemically and biologically inert and photo-stable. 4- Inexpensive. 5- Nontoxic. 1.5.1.1. Photocatalytic degradation The photocatalytic degradation process has gained a significant importance in the area of wastewater treatment. This process is a part of AOP which has proven to be a promising technology for the degradation of organic compounds. In comparison to other AOPs techniques, photocatalytic degradation is more effective because[39]: 1- Semiconductors are inexpensive. 2- Complete mineralization of various organic compounds. 3- Mild temperature and pressure is required. 4- No waste disposal problems. 1.5.1.2. Photocatalytic degradation of dyes The photocatalytic degradation of dyes is believed to occur according to the following mechanism[25]:

15


When the catalyst (semiconductor) is irradiated with visible light or UV radiations, electrons are promoted from the valence band to the conduction band, and an electron-hole pair is produced. Catalyst + hυ �⎯⎯⎯⎯⎯� e–cb + h+vb

1-19

Where, e–cb and h+vb are the electrons in the conduction band and the electron vacancy in the valence band, respectively. In most cases h+vb reacts with surface bound H2O to produce ●OH radicals, while, e–cb reacts with O2 to produce superoxide radical anion of oxygen. H2 O + h+vb �⎯⎯⎯� ●OH + H+ O2 + e–cb �⎯⎯⎯� O–2

●

1-20 1-21

●

The produced ●OH and O–2 can react with the dye to form other species, so the discoloration of the dye takes place. 1 2

O–2 ● + H2 O �⎯⎯⎯⎯� H2 O2

1-22

H2 O2 �⎯⎯⎯⎯� 2●OH

1-23

dye + e–cb �⎯⎯⎯⎯� dyered

1-25

●

OH + dye �⎯⎯⎯⎯� dyeox

1-24

Figure 1-6 depicts a mechanism for the generation of oxidative species in photocatalytic degradation technique.

16


Figure 1-6: Schematic representation for the mechanisms of generation of oxidative species in photocatalytic degradation technique[40].

1.5.2. Quantum yield Each process starting with absorption of a photon and ending with the disappearance of the molecule or its deactivation to the ground state is called a primary process. The quantum yield (Φi) of a primary process is defined as:

Φi =

Number of molecules undergoing that process Number of photons absorbed by the reactant

1-26

with both quantities referring to the same time interval[34].

1.6. Semiconductors Solid-state materials are usually described by the Band theory, which has a close parallelism with Molecular Orbital theory. Molecules are characterized by HOMO and LUMO orbitals, as depicted in Figure 1-7. Clusters can be described by the combination of atomic orbitals that

17


gives rise to closely spaced filled orbitals in addition to a set of closely spaced empty orbitals at higher energy[34].

LUMO

HOMO

Figure 1-7: Electronic energy levels in (a) a molecule, (b) a cluster, (c) a quantum dot and (d) a bulk semiconductor[34]. A solid may be thought of as consisting of a large number, say N, of atoms initially separated from one another. At relatively large separation distances, each atom is independent of all the others and the atomic energy levels and electron configuration appears as if isolated. However, when atoms come close to one another, electrons are perturbed by the electrons and nuclei of adjacent atoms, which lead to the splitting of each distinct atomic state into a series of closely spaced electron states in the solid to form what is known as electron energy band[41]. The band formed by the HOMOs is called valence band (VB), while the LUMOs form the conductance band (CB)[42]. Bandgap energy (Eg) is the energy separation between VB and CB. Depending on the value of Eg, solids may display conductor, semiconductor or insulator properties. The materials characterized by Eg in the range of 1‒ 4 eV are usually classified as semiconductors[34]. 18


Positions and width of energy bands of selected semiconductors are presented in Figure 1-8 and compared to those of TiO2.

Figure 1-8: Position and width of the energy band of TiO2 and several other illuminated semiconductors with respect to the electrochemical scale (NHE: normal hydrogen electrode)[43].

Bandgaps of the semiconductors are divided into two types: i. Direct bandgap: When the electron is excited from the valence band to the conductance band with no change in its momentum ∆k = 0, the semiconductor is said to have a direct band gap, and the absorption and emission of light are more effectual. GaS, ZnO and CdTe are examples of semiconductors with a direct band gap. ii. Indirect bandgap: The band gap of a semiconductor is indirect when the k at the valance band maximum differs from the k at the conduction band minimum. The transition of electrons in such a case is forbidden, hence, the momentum cannot be conserved. For the photon to be absorbed, coupling to a lattice vibration (a phonon) is required to compensate for the change in the wave vector during the transition. GaP, TiO2 and CdS are examples of indirect band gap semiconductors.

19


Figure 1-9: A sketch of the energy bands for (a) a semiconductor with a direct bandgap and (b) a semiconductor with an indirect bandgap[44]. Semiconductors occur in many different chemical compositions with a large variety of crystal structures. They can be elemental semiconductors, such as Si, carbon in the form of C60 or nanotubes and selenium (Se), or binary compounds such as gallium arsenide (GaAs), or organic compounds such as polyacetylene (CH)n, or oxides like CuO, ZnO, TiO2, or layered such as PbI2, MoS2 and GaSe, or magnetic like EuS and alloys such as Cdx–1MnxTe[45].

1.6.1. Classification of semiconductors Semiconductors are classified as depicted in Figure 1-10.

20


Semiconductors Extrinsic (impure)

Intrinsic (pure)

N-Type

P-Type

Figure 1-10: Schematic description of classification of semiconductors. 1.6.1.1. Intrinsic semiconductors Intrinsic semiconductors are those in which the electrical behavior is based on the electronic structure inherent to the pure material, and are characterized by a completely filled valence band separated from an empty conduction band by a band gap generally less than 2 eV. Silicon (Si) and Germanium (Ge) are elemental semiconductors with band gap energies 1.1 and 0.7 eV, respectively. In addition, a large number of compound semiconductors also show intrinsic behavior such as gallium arsenide (GaAs), indium antimonide (InSb), cadmium sulfide (CdS) and zinc telluride (ZnTe)[41]. 1.6.1.2. Extrinsic semiconductors The electrical characteristics of the extrinsic semiconductors are determined by impurity atoms, and virtually all commercial semiconductors are extrinsic[41]. Extrinsic semiconductors are classified into two types: n-type and p-type. Table 1-6 shows a comparison between these two types.

21


Table 1-6: Comparison between p-type and n-type semi-conductors [41, 46]. P-Type semiconductor 1. Formed when trivalent impurity like boron and lithium is added to a pure semiconductor.

N-Type semiconductor 1. Formed when pentavalent impurity like phosphorous and antimony is added to a pure semiconductor

2. The added impurity is known as the 2. The added impurity is known as the acceptor. donor. 3. There is a deficiency in electrons and excess of holes.

3. There is a deficiency in holes and excess of electrons.

4. The majority carriers are holes while minority carriers are electrons.

4. The majority carriers are electrons while minority carriers are holes.

5. The charge is conducted via holes.

5. The charge is conducted via electrons.

6. Fermi level is near valance band.

6. Fermi level is near conduction band.

1.7. Titanium Dioxide TiO2 also known as titanium white[47], has three polymorphs found in nature, namely anatase, rutile and brookite (note Figure 1-11)[48]. Formerly known as octahedrite, anatase is a polymorph of titanium dioxide, and its name comes from the Greek word anatasis, which means “extension”- a

22


reference to the elongate octahedral crystals that are the most common habit of anatase[49]. Anatase is the strongest oxidizer and exhibits the highest photocatalytic activity among the three polymorphs, whereas the presence of rutile with anatase only enhances its photocatalytic activity[50]. Anatase is the less thermodynamically stable polymorph of TiO2, although energy calculations demonstrate that this phase appears as the more likely phase when the grain size is around 10 nm[38].

Figure 1-11:Unit cells of TiO2 rutile(a), anatase(b) and brookite(c)[51]. Rutile takes its name from the Latin rutilis, which means “red” or “glowing”[49]. Rutile is best known as white pigment, and this commercial application arises from the fact that fine particles scatter incident light extremely strongly; even crystals of TiO2 possess a very high refractive index (μ = 2.6 for rutile, 2.55 for anatase)[52], hence it is used in sunscreen lotions to protect skin from UVrays[48] and it is particularly useful in cosmetics, and it is an important semiconductor for optical instruments and polarized optics and has several applications such as paints and plastics[50]. 23


Pure rutile is produced by converting TiO2 ore into TiCl4 by treating it with Cl2 and C at 1200 K then oxidizing the product by O2 at ≈ 1500 K[52]. Brookite is named in 1825 after British crystallographer H.J. Brooke[49]. Its photocatalytic activity has been not much studied due to the difficulties in the preparation of pure brookite without rutile and anatase. In recent years, the interest in brookite has increased and pure brookite has proved to be an interesting candidate in photocatalytic applications[53]. Table 1-7 summarizes some of the crystal structure data for the different polymorphs of TiO2.

1.7.1. Photolysis effect: Since the discovery of water photolysis effect of TiO2 by Fujishima and Honda in 1972[54], it has been paid much attention and used widely in photocatalytic degradation of pollutants, photocatalytic CO2 reduction into energy fuels, splitting of water, solar cells, biomedical devices and lithium ion batteries[55]. TiO2 is not activated by visible light, but by ultraviolet light[56], and this is due to its wide band gap (see Table 1-7) which leads to absorption of light with wavelengths only below ∼390 nm[57]. Nevertheless, it is advantageous over other semiconductors because it is chemically and biologically inert, photocatalytically stable, easy to produce and to use, catalyzes reactions efficiently, cheap and with no risks to environment or humans[56]. A wide range of pollutants can be successfully destroyed using TiO2 as a catalyst, including hydrocarbons and halogenated organic compounds as well as some herbicides, pesticides and dyes[52].

24


Table 1-7: Crystal structure data for TiO2 [58, 59] Properties

Rutile

Anatase

Brookite

Crystal structure

Tetragonal

Tetragonal

Orthorhombic

a = 4.5936

a = 3.784

a = 9.184

c = 2.9587

c = 9.515

b = 5.447

Crystalline structure

Lattice constant (Å)

c = 5.154 Molecule (cell)

2

2

4

Volume/molecule

32.2160

34.061

32.172

Density (g cm-3)

4.13

3.79

3.99

Ti–O bond length (Å)

1.949 (4)

1.937 (4)

1.87 – 2.04

1.980 (2)

1.965 (2)

81.2°

77.7°

90.0°

92.6°

2.98

3.05

3.26

7.0 – 7.5

5.5 – 6.0

5.5 – 6.0

(Å3)

O–Ti–O bond angle

Experimental band

77.0° – 105°

gap (eV), at pH = 7 Mohs hardness TiO2 crystals in nature

25


1.7.2. TiO2 nanoparticles In recent years, TiO2 has received wide attention and has a good prospect in solar energy cells and in environmental purification. Recently, TiO2 has been utilized as a photocatalyst for various fields of environmental purification, such as decolorization of dyes, degradation of hazardous volatile organic compounds and purification of air and water[60]. Nano-TiO2 offers advantages due to its band-gap in the desired UV-visible spectral range as a photocatalyst. Besides its optical properties, TiO2 has been found to be a non-toxic and stable material, which is available at a lower cost[50]. TiO2 has been prepared in the form of powders, crystals, thin films, nanotubes and nanorods[58], and many synthesis methods, as shown in Figure 1-12, have been used to synthesize nanostructured TiO2 with different morphologies.

Sol-Gel

Vapor depostion

Slovothermal Electrochemical

Plasma evaporation

TiO2 Nanoparticles

Microwave

Chemical Vapor Condensation

Solution Combustion

Microemulsion

Gas-phase hydrolysis Sonochemical

Micelle and Inverse Micelle

Figure 1-12: Synthesis methods of TiO2 nanostructures [43, 61, 62].

26


1.7.3. Sol-gel method Sol-gel method has been used to prepare various kinds of metal oxides thin films, mainly because of their low cost and flexible applicability of wide ranges of size and shape of substrates[63]. These methods have diverse applications such as the development of new catalysts, chemical sensors and fibers, in addition to its usage in solid state electrochemical

devices,

besides

other

photochromic

and

nonlinear

applications[64]. The precursors of a typical sol-gel process are usually inorganic metal salts or metal alkoxides that undergo different types of hydrolysis and polycondensation reactions[65] as shown in Figure 1-13. TiCl4 + ROH

Precursor (Alkoxides)

Ti(OR)4 + 4HCl

Hydrolysis

Volatile Organic Compounds

Sol or colloidal suspension Condensation and Polymerization

Heat treatment

Gel Phase Calcination

TiO2 Nanostructures

TiO2 nanoparticles

Figure 1-13: Schematic diagram for Sol-gel method. Sol-gel method is a “bottom-up” approach which involves the connection of molecular building blocks such as SiO4 tetrahedra and TiO6 octahedra with each other step by step[66].

27


The chemistry of sol-gel methods involve the preparation of inorganic polymers or ceramics from solution through a transformation from liquid precursors to a “sol”, which is a stable suspension of colloidal particles (nanoparticles) in a liquid, and finally to a “gel”[67], which is an interconnected rigid network with pores of sub-micrometer dimensions and polymeric chains whose average length is greater than a micrometer[68]. In most cases, gelation occurs due to the formation of covalent bonds between the sol particles, and it can be reversed when other bonds are involved, i.e. van der Walls forces or hydrogen bonds[66]. The key steps in the sol-gel process can be summarized as follows[67]: 1- “Sol” is synthesized from hydrolysis and partial condensation of alkoxides. 2- “Gel” is formed through polycondensation to form either oxo M–O–M or hydroxo M–OH–M bridges between the metallic atoms M of the precursor molecules. 3- The gel is aged by continuous condensation resulting in shrinking and solvent expulsion. 4- The gel is dried to form a dense “xerogel” via collapsing of the porous network. The surface M–OH groups are removed via calcination at high temperature (up to 800°C). There are many advantages of sol-gel method[65, 69]: 1- It not only allows for materials to have any oxide composition, but it also permits the productions of new hybrid organic-inorganic materials not present in nature. 2- The products obtained by this method are very pure, and this purity can be obtained by purifying the precursors by distillation, crystallization or electrolysis. 28


3- The chemical processes of the initial steps are always carried out at ambient temperature; this minimizes the chemical interactions between the material and container walls. 4- The kinetics of the various chemical reactions can be easily controlled by the low processing temperature and the often-dilute conditions. 5- Particle size, shape and properties can be controlled finely.

1.8. Adsorption Adsorption is a surface phenomenon which involves the concentration of molecules of a gas or a liquid (adsorbate) at the surface of a solid material (adsorbent)[33]. Adsorption and desorption (the reverse process of adsorption) play an important role in heterogeneous photocatalytic processes used in pollutant degradation, as these processes involve the adsorption of pollutants on the surface of the photocatalyst which facilitates the conversion of the pollutants into CO2 and H2O[70].

1.8.1. Adsorption on semiconductors surface: The overall process of heterogeneous photocatalysis can be divided into five independent steps[71]: 1- Transfer of the reactants in the fluid phase to the surface. 2- Adsorption of one molecule of reactants at least. 3- Reaction in the adsorbed phase. 4- Desorption of the products. 5- Removal of the products from the interface region. Steps (1) and (5) depend on the amount and the particle size of the photocatalyst and on the concentration of each of the reactants and products.

29


Steps (2) to (4) depend on the competition of the reactants and products molecules on the active sites of the photocatalyst. Step (3) is responsible for the production of the reactive radical (mainly OH radical) which is regarded the first step in the chemistry of AOPs. The mechanism of the photocatalytic oxidation by semiconductors like TiO2 consists the following processes[72]: 1- (e– - h+) pairs in the semiconductor particles are generated: TiO2 + hv

(h+ – e–) (exciton)

h + + e–

1-27

This process involves the excitation of an electron from the valence band to the conduction band via the absorption of light with energy equal to or greater than the bandgap of the semiconductor. 2- At the surface, the semiconductor can donate electrons to reduce an electron acceptor, similarly, the holes can migrate to the surface and combine with an electron from a donor species and oxidize it. The trapping of the charge carrier proceeds by the following mechanism: (100 ps)

1-28

(10 ns)

1-29

3- Recombination of the electron and the hole can take place in the volume of the semiconductor particle or on the surface. This process is accompanied by the release of heat.

30

(100 ns)

1-30

(10 ns)

1-31


1.8.2. Adsorption of oxygen The reactive species in an n-type semiconductor is the electron, and in order to keep the photocatalytic degradation process going the accumulation of the electrons on the particles must be avoided, since it may lead to an increase in the rate of hole-electron recombination rate and lower the quantum yield. Hence, the presence of an electron acceptor in the solution is essential for the efficiency of the photocatalytic degradation process. Usually, O2 is the most common electron acceptor used in such processes, due to its availability at low or no cost and its ability to dissolve in aqueous and other solutions[73]. The photo-generated electrons react with molecular oxygen to form active species like superoxide anions (O•– 2 ) and hydrogen peroxide (H2O2)[74]. Figure 1-14 depicts the process of adsorption of oxygen molecules on the surface of TiO2 and the production of H2O2 molecules.

Figure 1-14: Schematic description of adsorption of oxygen on the surface of TiO2[75]. The adsorption process of O2 gas on the surface of TiO2 and the formation of different species can be expressed as follows[40, 76]: O 2gas  → O 2ads

31

1-32


→ O-2ads O2ads + e- 

1-33

→ 2O-ads O-2ads + e- 

1-34

→ O 2-ads O-ads + e- 

1-35

O-2ads +H +  →  OOH

1-36

→ H 2O2 + O2 O-2ads +  OOH + H + 

1-37

→ 2  OH H 2O2 

1-38

1.8.3. Adsorption of dyes The adsorbed dyes molecules on the surface of the semiconductor can absorb a radiation in the visible range besides the radiation with a short wavelength and with energy which is less than the Eg of the radiated semiconductor. The dye molecules will suffer decolorization through a free radical mechanism which is summarized by the following equations[40]:

1

Dye + hν (Vis or UV) �⎯⎯⎯�

1

3

Dye* or Dye*

Dye● or Dye● + Semiconductor �⎯⎯⎯⎯� Dye●+ + e–semiconductor 3

e–semiconductor + O2 �⎯⎯⎯⎯� O●– 2 + Semiconductor

Dye●+ + O●– 2 �⎯⎯⎯⎯� DyeO2 �⎯⎯⎯⎯� degradation products Dye●+ + HO●2 (or HO● ) �⎯⎯⎯⎯� degradation products Dye + 2HO● �⎯⎯⎯⎯� H2 O + oxidation products Dye●+ + HO- �⎯⎯⎯⎯� Dye + HO●

Dye●+ + H2 O �⎯⎯⎯⎯� Dye + HO● + H+

32

1-39 1-40 1-41 1-42 1-43 1-44 1-45 1-46


1.9. Dyes Dyes can be classified according to chemical structure or the mode of application. Table 1-8 and depict in Figure 1-15 the different types of dyes classification. Table 1-8: Classification of dyes according to chemical structure [77, 78]. Nitroso

Indamine

Nitro Azo Azoic Stilbene Carotenoid Diphenylmethane Triarylmethane Xanthene Acridine Quinioline Methine Thiazole Oixidation bases

Indophenol Azine Oxazine Thiazine Sulfur Lactone Aminoketone Hydroxyketone Anthraquinone Indigoid Phthalocyanine Natural organic Inorganic pigments

Reactive Vat

Disperse Dyes Acid

Direct Basic

Figure 1-15: Classification of dyes according to application method[77].

33


1.9.1. Triarylmethane dyes: Triarylmethane dyes (known also as Triphenylmethane dyes) belong to the class of polymethine dyes[79]. In this class of dyes two aryl rings and R group bonded to the central methine carbon create the branches, in which the polymethine chain is combined. The R group contains π-electrons or lone pairs of electrons capable of interaction with the π-electron system[80]. The chromophores of triarylmethane dyes are the quinonoid group, and the different types of dyes are obtained by the introduction of NH2, NR2 or OH groups into the para positions of Benzene ring of triarylmethane[81], and only di- and trisubstituted derivatives are of commercial value[82]. The triphenylmethane dyes are among the oldest industrially product dyes. They have a brilliant color because of resonance, and cover a range of shades from red to blue including violet and green. At first they were prepared by methods discovered empirically, without knowledge of their constitution. The first industrial production process for Fuchsine (see Figure 1-16) was developed by Verguin in 1859[80].

Figure 1-16: The chemical structure of Fuchsine. Triarylmethane dyes do not show a good resistance to light, however, they are one of the most important groups of synthetic dyes, because they are brilliant and have tinctorial strength, and low cost[83].

34


1.9.2. Light Green dye: Light Green SF Yellowish (LGSF), also known as Lissamine green SF is an triarylmethane-type dye (see

Figure 1-17) with the structural formula

C37H34N2Na2O9S3 and a molecular weight of 792.85 g.mol-1[84]. SO3Na CH3 N

-

SO3Na

N

O3S

CH3

Figure 1-17: Chemical structure of LGSF[84]. LGSF is used as a stain in cell, cytoplasm, endoscope, microorganisms, eye membranes, retina, proteins and hairs. It also has several biological applications and is used in cosmetics, oral hygiene products, sunscreen, detecting proteins and treating apolipoprotein E-related diseases. In addition, LGSF has multiple industrial applications and is used in color filters, recording materials, inks, highlighters, adhesives, photographic materials, detergents, textiles and leathers[85]. LGSF can severely affect the metabolic system, and it has the ability to accumulate and permeate in the case of skin contact and it acts as irritant in the case of ingestion or inhalation. In addition, the dye and its metabolites can induce carcinogenic effects in living systems. Schiller reported that LGSF can produce sarcomas and overexposure to this dye may lead to a type of blood disorder called methemoglobinemia. Allmark and his coworkers reported also the chronic toxicity of this dye[86]. 35


1.10. Aims of the study This work is composed of three main parts: 1. The first part investigates the synthesis of TiO2 nanoparticles via sol-gel method. 2. The second part includes characterizations of the obtained samples employing XRD, AFM and Fluorescence spectroscopy analysis, in order to compare the prepared samples with commercial and commercial nano TiO2. 3. The third part investigates the effects of various parameters on the photocatalytic degradation of LGSF dye, such as: a. Dose of catalyst. b. Dye concentration. c. Initial pH of solution d. Temperature of solution. The results are used to estimate the best (optimum) conditions for degradation of LGSF dye.

36

Chapter Two

Experimental Part

Chapter two: Experimental part

2.1. Chemicals Table 2-1 lists the chemicals used in this work. All the chemicals were obtained and used without further purification.

Table 2-1: Chemicals and their commercial sources. No. Chemicals

Purity % Supplied from 99

Avantor Performance Materials, Inc.

1

Light Green SF Yellowish

2

Titanium tetrachloride (TiCl4)

> 99

BDH, England

3

Absolute ethanol (C2H5OH)

99.9

J.T.Baker, Netherlands

4

Hydrochloric Acid (HCl)

37

Merck, Germany

5

Sodium hydroxide (NaOH)

99

Sigma Aldrich

6

1,10-Phenanthroline

99

Riedel-De-Haen AG, Germany.

7

Iron(II) sulfate (FeSO4)

99

Fluka, Switzerland.

8

Sulfuric acid

98

9

Potassium oxalate (K2C2O4)

10

Titanium(II) oxide (nanoparticles)

99

11

Titanium(II) oxide

99

98.5

37

Himedia Chemical Company Riedel-De-Haen AG, Seelze, Hannover, Germany Hefei EV NANO Technology Co., Ltd., China Fluka, Germany


2.2. Instruments Table 2-2 lists types of instruments and the supplier companies. Table 2-2: Instruments used in this work. No. 1

Instrument

Company

Place

Hot plate magnetic LMS1003/ Labtech/ Daihan University of Karbala stirrer

lab Techco, LTD

2

Oven

Memmert, Germany.

University of Karbala

3

Spectrophotometer

Spectro SC, LaboMed, Inc.


4

UV-Vis

Cary 100Bio, Shimadzu, University of Karbala

Spectrophotometer

Japan

Fluorescence

Scinco FS-2, Korea

University of Babylon

Force Angstrom AA-3000, USA

University of Baghdad

5

Spectrometer 6

7

Atomic Microscope

University of Baghdad

X-Ray Diffraction

Angstrom DX-2700, USA

University of Babylon

University of Babylon 8

High

Pressure Rudium, China


Mercury Lamp UV A (400 W) 9

pH meter

Hanna

Instruments, University of Karbala

Mauritius 10

Sensitive Balance

BL

210

S,

Sartorius- University of Karbala

Germany 11

Centrifuge

Hettich EBA20 – Germany


12

Ultrasonic bath

FALC, Italy


38


2.3. Preparation of TiO2 nanoparticles by sol-gel method TiO2 nanoparticles had been synthesized by sol-gel method, which was developed from a previously described procedure[60], as follows: 10, 25 and 30 mL of TiCl4 were introduced dropwise into 100 ml of absolute ethanol at ambient temperature with continuous stirring to obtain three mixtures of volume ratio 1:10, 1:4 and 3:10 respectively. TiCl4 is a highly volatile compound, thus the reactions were performed under fume hood. HCl gas was exhausted and a light-yellow solution was obtained from each mixture. It was gelatinized initially under air atmosphere and inside the hood using a heater at 80°C until a white gel appeared. The obtained gel was dried at 85°C in an oven for 20 hours and a sol-gel was formed. Finally, the sol-gel was grinded and divided and calcinated in furnace for 2 hours at 600 and 800°C. According to each experiment, different phases of TiO2 nanopowder were obtained. Figure 2-1, to Figure 2-3 illustrate this process.

Figure 2-1: Synthesis setup for TiO2 nanoparticle preparation using sol-gel method.

39


TiCl4 Dropwise under continuous stirrer 10 ml

25 ml

30 ml

C2H5OH 100 ml

C2H5OH 100 ml

C2H5OH 100 ml

Vol% =

1:10

1:4

3:10

Yellow suspension solutions Drying at 85 °C in an oven for 15h

TiO2 Powder Grinding, Dividing heat treatment for 2h at 600 °C

800 °C

TiO2 Nanostructures

Figure 2-2: Detailed schematic representation for experimental procedure of synthesis of TiO2 nanoparticles using sol-gel method.

40


TiCl4 + 4C2H5OH

Ti(OC2H5)4 + 4HCl

Volatile Organic Compounds

Heat treatment 600 and 800 °C

TiO2 Nanostructures Figure 2-3: Schematic diagram for experimental procedure

2.4. X-Ray Diffraction patterns The characterization of the crystal morphology and size of the commercial and prepared TiO2 was carried out via XRD measurements. The average crystallite size, in nm, was determined by Debye-Scherrer formula[87]:

L=

kλ β cosθ

2-1

where L is the average crystallite size, k = 0.90 is the constant crystal lattice, λ = 0.154 nm which is the wavelength of the radiation, β is the full width at half maximum in radians, θ is the position of the maximum diffraction. In addition, modified Scherrer equation was employed to calculate the crystallite size (L`) more accurately.

41


= ln β ln Plotting ln β against ln

kλ 1 + ln L` cosθ

2-2

1 , a straight line, which is theoretically has a slope cosθ

equal to 1, is obtained. Crystallite size may be calculated by getting the intercept and obtaining its exponential:

e

ln k λ L`

=kλ L`

2-3

where k is substituted with 0.9 and λ with 0.154 nm.

2.5. Atomic Force Microscopy (AFM) The AFM images were recorded with (SPM-AA3000) instrument after a known amount of each sample was suspended in ethanol and treated by ultrasonic instrument for 10 min in power of 25 kHz. A drop of each of the obtained colloidal solutions was deposited on (1 × 3 cm) glass slides.

2.6. Band gap Energy measurement Fluorescence spectroscopy was used to calculate the band gap energy for the commercial and prepared TiO2 nanoparticle, employing Tauc equation[88]:

Band gap (eV) =

1240 λ (nm)

2-4

Where λ is the wavelength that corresponds to the maximum intensity in the fluorescence spectrum.

42


2.7. Light Intensity measurement The light intensity was calculated via calculating the light flux density using the actinometric method. The most commonly utilized method in actinometry is based on a chemical actinometer which undergo a photochemical reaction with a known quantum yield[89]. The ferrioxalate actinometric solution was prepared in the same photocatalytic reactor with similar volume of the reaction mixture (100 mL) by mixing 40 mL of 0.15 M of Fe2(SO4)3, 50 mL of 0.45 M of K2C2O4 and 10 mL of 0.05 M of H2SO4. The mixture was irradiated under atmospheric oxygen using mercury lamp UV (A) for 15 min. The color of the mixture was changed to yellowish green indicating the production of K3[Fe(C2O4)3].3H2O as shown in Figure 2-4.

Figure 2-4: The experiment of light intensity measurement. In a regular interval (5, 10 and 15 min) 2.5mL of the irradiated solution was drawn and centrifuged (3000 rpm, for 10 min). After filtration, 0.5 mL of 1,10phenonethroline was added to the solution and a red orange complex was formed. Under light excitation the potassium ferrioxalate decomposes according to the following equations[90]: 43


hν

•– 2– 2+ Fe(C2 O4 )3– 3 �⎯⎯⎯⎯⎯� Fe + C2 O4 + 2C2 O4 ∆

•– 2– 2+ Fe(C2 O4 )3– 3 + C2 O4 �⎯⎯⎯⎯⎯� Fe + 2CO2 + 3C2 O4

2-5 2-6

The quantity of ferrous ions formed during the irradiation period is monitored by conversion to the colored tris-phenanthroline. The light intensity had been calculated via following equations: I0 =

I0 =

A V 1V 3 ε t Qλ V 2

1.768 ×100 × 3 10453×103 × 600 ×1.2 × 0.5

I0 =1.409 × 10–7 Einstein s–1

2-7

2-8 2-9

where: I0 is light intensity, A is absorbent at 510 nm, V1 is volume of solution irradiated (100 cm3), V3 is the final volume after complexation with 1,10phenonethroline (3 cm3), t is irradiation time, Q is quantum yield (1.2)[89], V2 is the aliquot of the irradiated solution taken for the determination of the ferrous ions (0.5 cm3) and ε is molar absorbent coefficient (slope value that calculated from calibration curve of Fe(II)). So, the calibration curve of Fe(II) was obtained by using a series of solutions prepared from a stock solution (1×10–3 M) of FeSO4.7H2O dissolved in 0.1 M of H2SO4. Then 0.5 mL from each of these solutions was added to 2.5 mL of 0.1% of 1,10-phenanthroline. The absorbance of the resultant tris-phenantrholine complex was measured at 510 nm using a UV-visible spectrophotometer. The molar absorptivity ε of the complex at 510 nm was found to be 1.0453 × 104 L mol–1 cm–1. Typical calibration values are given in Table 2-3 and plotted in Figure 2-5.

44


Table 2-3: Absorbance at different concentrations of Fe(II). [Fe(II)] / 104 M

Abs in 510 nm

0.200

0.218

0.400

0.403

0.600

0.609

0.800

0.828

1.000

1.067

1.2 R² = 0.9974

1

Abs

0.8 0.6 0.4 0.2 0

0

0.2

0.4

0.6

0.8

1

1.2

[Fe(II)] / 104 M Figure 2-5: Calibration curve for Fe(II) as a complex with 1,10phenanthroline.

2.8. Measurement of maximum absorption wavelength of the dye UV-Visible spectroscopy of the diluted LGSF dye aqueous solution was recorded to investigate λmax of the dye. The spectrum is shown in Figure 2-6. The UV-Visible spectrum shows that LGSF dye has two absorption peaks, 233 and 640 nm.

45


Figure 2-6: UV-Visible spectrum of LGSF dye aqueous solution.

2.9. Calibration curve of the dye The calibration curve was obtained by using standard LGSF dye aqueous solutions. The absorbance of each concentration was measured at 640 nm. Typical calibration values are given in Table 2-4 and plotted in Figure 2-7. Table 2-4: Absorbance at different concentrations of LGSF. [LGSF] (ppm)

Abs in 640 nm

1 3 5 7 10 13 15 18 20

0.156 0.462 0.551 0.656 0.896 1.135 1.259 1.419 1.516

46


2

R² = 0.9599

Abs

1.5 1 0.5 0

0

5

10

15

20

25

[LGSF] ppm Figure 2-7: Calibration curve at different concentration of LGSF Dye

2.10. Measurement of the optimal conditions The photocatalytic activity experiments were done using a 400 mL photoreactor, and a high-pressure mercury lamp UV(A) with 400 W was used as a source of radiation (see Figure 2-8). Prior to irradiation, suspensions were stirred in darkness for 30 min using magnetic stirrer to ensure adsorption equilibrium. During adsorption and irradiation, suspensions were sampled at regular intervals. 4 cm3 of the reaction mixture was collected and centrifuged for 15 min. The supernatant was carefully removed by a syringe with a long pliable needle and centrifuged again at same speed period of time. The second centrifuge was found necessary to remove fine particle of the catalyst. After the second centrifuge, the absorbance at the maximum wavelength of LGSF (640 nm) was measured with UV-visible spectrophotometer.

47


Figure 2-8: The photocatalytic reactor.

2.10.1. Optimal concentration The optimal concentration of the dye was measured by adding 175mg of the catalyst into a series of Light Green SF Yellowish solutions of (15, 20, and 25 ppm in 100 mL) with TiO2 nanopowder, and (25, 30, 40 and 50 ppm in 100 mL) with commercial Degussa P-25.

2.10.2. Optimal catalyst dose To measure the optimal amount of the catalyst two sets of experiments were carried out. The first set was carried out using (50, 100, 150, 175, 200, 250, 300 and 400 mg) of TiO2 nanopowder introduced into a dye solution of 20 ppm in 100mL at ambient temperature. The second was carried out using (50, 100, 200, 250, 300 and 400 mg) of commercial Degussa P-25 which was introduced into a dye solution of 30 ppm in 100 mL at ambient temperature.

48


2.10.3. Optimal pH The measurement of the optimal pH for each type of the catalyst was done through a set of experiments with pH of 3, 4, 5, 6, 7.3 and 8 respectively.

2.10.4. Optimal temperature To measure the optimal temperature a set of experiments for each type of the catalyst was carried out at (10, 15, 20, 25, 30 and 35 °C) respectively. In each experiment the system was stirred for 30 minutes to reach the adsorption desorption equilibrium, then irradiated with UV light (Rudium, 230 volts, 50Hz, 400w, UV A) at ambient temperature. Then, 4mL of the suspension was extracted from the reactor in a regular interval of 5 minutes and centrifuged twice at a rate of 6000 rpm to remove the catalyst. The second centrifuge was found necessary to remove fine particle of the catalyst. The absorbency of the supernatant solution was measured using a UV-Vis spectrometer at the maximum absorption wavelength (640 nm) of LGSF.

2.11. Photocatalytic kinetic analysis According to Langmuir-Hinshelwood (L-H) model, with the assumption of the initial dye concentration to be low; the reaction follows the pseudo first-order kinetics, so the apparent rate constant (kapp) expression was calculated using the following equations[91]:

C t = C 0e

( − k app . t )

2-10

Where: C0: is the initial concentration of LGSF dye at irradiation time equal to zero minutes. Ct: is the concentration of the same dye at time (t) of irradiation.

49


Ct (−k . t ) = e app C0

2-11

Ct = − k app . t C0

2-12

Or

ln

ln

C0 = k app . t Ct

The Langmuir-Hinshelwood (L-H) kinetic model has four possibilities[91–93]: 1- The reaction occurs between two adsorbed substances (Dyeads and •OHads): •

OHads + Dyeads �⎯⎯⎯⎯⎯⎯� Decolorization or degradation

2-13

2- A nonbound radical (radical in solution) reacts with and adsorbed dye molecule (Dyeads and •OHsol): •

OHsol + Dyeads �⎯⎯⎯⎯⎯⎯� Decolorization or degradation

2-14

3- The reaction occurs between a radical linked to the surface and substance molecules in solution (Dyesol and •OHads): •

OHads + Dyesol �⎯⎯⎯⎯⎯⎯� Decolorization or degradation

2-15

OHaqu + Dyesol �⎯⎯⎯⎯⎯⎯� Decolorization or degradation

2-16

4- The reaction results from two free species in solution (Dyesol and •OHaqu): •

In all cases, the equation of the reaction rate is expressed in a similar manner.

2.12. Photocatalytic decolorization efficiency The photocatalytic decolorization efficiency (PDE) was calculated using the following equation[94]: PDE =

(C 0 − C t ) ×100 C0

50

2-17


2.13. Activation Energy Arrhenius equation was employed to calculate the activation energy in the range of temperatures (283.15 – 308.15 K):

E lnk app = − a + ln A RT

2-18

where: kapp is apparent rate constant, T is temperature of reaction, Ea is the apparent activation energy, R is the gas constant (8.314 J mol–1 K–1), and A is frequency constant.

2.14. Thermodynamic Parameters Eyring–Polanyi equation was employed to calculate ∆H‡ and ∆S‡ values[95]: ‡ k  ∆H ‡ 1  k B  ∆S ln + = − +    R T R T   h 

ln 

2-19

where k is the apparent rate constant, T is the absolute temperature, ∆H‡ is the enthalpy of activation, R is the gas constant, kB is Boltzmann’s constant, h is Plank's constant and ∆S‡ is the entropy of activation. The free energy ∆G‡ of activation was calculated via equation 2-20: ∆G‡ = ∆H‡ – T∆S‡

51

2-20

Chapter Three

Results

Chapter Three: Results

3.1. Catalysts characterization 3.1.1. X-Ray Diffraction patterns The XRD patterns of TiO2 nanoparticles obtained via sol-gel method in different volumetric ratios of addition and different annealing temperatures are shown in Figure 3-1 to 3-8 respectively.

Figure 3-1: X-ray diffraction pattern for the commercial TiO2.

Figure 3-2: X-ray diffraction pattern for the commercial nano TiO2.

52


Figure 3-3: X-ray diffraction pattern for the prepared TiO2 nanoparticles with 600 °C calcination temperature. [V/V (TiCl4:EtOH) = 1:10].


53




54




55

Chapter Three: Results The average crystallite size (L) and the crystallite size (L`) were calculated employing Scherrer and modified Scherrer equations respectively. The results are depicted in Figure 3-9 and listed in Table 3-1. -5.2

y = -4.5249x - 5.1965 R² = 0.954

-5.2 -5.25 -5.3 -5.35 -5.4 -5.45 -5.5

-5.45 -5.5 -5.55 -5.6 -5.65 -5.7

0

0.05 0.1 ln (1/cosθ)

0.1

y = -0.1709x - 5.5793 R² = 0.0088

e

0

0.05

0.1

y = 2.2373x - 5.5031 R² = 0.4014

0.05

0.1

0

0.05 0.1 ln (1/cosθ)

0.15

f

-5.5 -5.6

g

0.15

ln (1/cosθ)

y = -1.9985x - 5.3923 R² = 0.3831

-5.7 -5.8

0.15

-5.4 0

d

-5.4

-5.3 -5.5

0.1

-5.7 -5.8

ln β

ln β

-5.2

0.05 ln (1/cosθ) y = 0.813x - 5.6917 R² = 0.0657

-5.6

ln (1/cosθ) -5.1

0

-5.5

y = 0.8998x - 5.3765 R² = 0.0557 0.05 ln (1/cosθ)

y = -8.6022x - 5.0969 R² = 0.9907

-5.8

-5.4

c

0

-5.6 -6

0.15

ln β

ln β ln β

ln β

-5.6 -5.8

b

-5.4

ln β

ln β

-5.4

-5.2

a

-5 -5.1 -5.2 -5.3 -5.4 -5.5

0

0.05 0.1 ln (1/cosθ)

0.15

y = -4.8404x - 4.9785 R² = 0.9876

0

0.05 0.1 ln (1/cosθ)

h

0.15

Figure 3-9: Modified Scherrer equation of TiO2, a) prepared (1:10 at 800°C), b) prepared (3:10 at 800°C), c) prepared (1:4 at 800°C), d) prepared (1:10 at 600°C), e) prepared (3:10 at 600°C), f) prepared (1:4 at 600°C), g) commercial bulk, h) commercial nano. 56

Chapter Three: Results Table 3-1: Crystallite size and average crystallite size of commercial bulk, commercial nano and prepared TiO2. Sample of TiO2

Crystallite size

Average crystallite

(L`) / nm

size (L) / nm

prepared (1:10 at 800°C)

25.0

38.1


22.7

48.1


30.0

31.5


41.1

48.0


36.7

44.8


30.5

39.7

Commercial bulk

34.0

33.8

Commercial nano

20.1

27.0

3.1.2. Atomic Force Microscopy (AFM) The two and three dimensions AFM graphs of the commercial bulk, commercial nano and prepared TiO2 in different volumetric ratios with calcination temperatures at 600 and 800 °C are shown in Figure 3-10 to Figure 3-17 respectively.

57


Figure 3-10: AFM images of the commercial TiO2, a) 2-dimensional image, b) 3-dimensinal image, c) Granulity cumulation distribution chart.

58


Figure 3-11: AFM images of the commercial nano TiO2, a) 2-dimensional image, b) 3-dimensinal image, c) Granulity cumulation distribution chart.

59


Figure 3-12: AFM images of the prepared TiO2 nanoparticles with 600°C calcination temperature. [V/V (TiCl4:EtOH) = 1:10], a) 2-dimensional image, b) 3-dimensinal image, c) Granulity cumulation distribution chart.

60


Figure 3-13: AFM images of the prepared TiO2 with 800°C calcination temperature. [V/V (TiCl4:EtOH) = 1:10], a) 2-dimensional image, b) 3-dimensinal image, c) Granulity cumulation distribution chart.

61


Figure 3-14: AFM images of the prepared TiO2 with 600°C calcination temperature. [V/V (TiCl4:EtOH) = 1:4], a) 2-dimensional image, b) 3dimensinal image, c) Granulity cumulation distribution chart.

62


Figure 3-15: AFM images of the prepared TiO2 with 800°C calcination temperature. [V/V (TiCl4:EtOH) = 1:4], a) 2-dimensional image, b) 3dimensinal image, c) Granulity cumulation distribution chart.

63



64



65

Chapter Three: Results Table 3-2: Particles size of commercial bulk, commercial nano and prepared TiO2 samples. Sample

Particle size/nm


100.16


98.96


100.20


163.91


96.77


109.1

Commercial bulk

432.93

Commercial nano

66.72

3.1.3. Band gap energy measurements Fluorescence spectra of commercial bulk, commercial nano and prepared TiO2 were recorded to investigate the optical band gap energy. The spectra are shown in Figure 3-18 to Figure 3-20 respectively and the results obtained via equation 25 are listed in Table 3-3. 140 120 Intensity

100 80 60 40 20 0

200

300

400

500

600

wavelength/nm

Figure 3-18: Fluorescence spectrum of commercial bulk TiO2.

66

700

800


350 300 250 Intensity

200 150 100 50 0

200

300

400

500

600

700

800

wavelength/nm

Figure 3-19: Fluorescence spectrum of commercial nano TiO2. 140 120 100 Intensity

80 60 40 20 0

200

300

400

500

600

700

800

wavelength/nm

Figure 3-20: Fluorescence spectrum of prepared TiO2 nanoparticles with 600°C calcination temperature [V/V (TiCl4:EtOH) = 3:10].

67

Chapter Three: Results Table 3-3: Band gap energy measurements results. Parameters

Commercial

Commercial

Prepared

TiO2

nano TiO2

TiO2

λ / nm

348.2

350.6

349.2

Eg / eV

3.56

3.53

3.55

3.2. Photocatalytic degradation of LGSF dye 3.2.1. Preliminary experiments A series of experiments were conducted under light intensity of 1.409 × 10–7 Einstein s–1 to obtain the essential conditions of the photocatalytic reaction. 3.2.1.1. Dark reaction Two reactions were carried out in the absence of UV light. The first reaction was carried out using 20ppm of LGSF dye with the addition of 0.175 g of commercial TiO2, while the second reaction was carried using the same solution but with addition of commercial nano TiO2. The results, which are expressed in Table A-1 and plotted in Figure 3-21, show that there is no reaction takes place in the absence of UV radiations.

68


ln(C0/Ct)

4.5E-03 4.0E-03 3.5E-03 3.0E-03 2.5E-03 2.0E-03 1.5E-03 1.0E-03 5.0E-04 0.0E+00

Commercial Commercial nano

R² = 0.9618

R² = 0.6786

0

10

20

t/min

30

40

50

60

70

Figure 3-21: The variation of ln(C0/Ct) with time in the absence of radiation. 3.2.1.2. Photolysis of LGSF dye The photolysis reaction for 20ppm of LGSF solution at 293 K was carried out in the absence of the catalyst. The results, which are expressed in and plotted in Figure 3-22, show that there is no reaction takes place in the absence of the

ln(C0/Ct)

catalyst. 0.12 0.1 0.08 0.06 0.04 0.02 0

R² = 0.9759

0

10

20

30

40

50

60

70

t/min

Figure 3-22: The variation of ln(C0/Ct) with time in the absence of the catalyst (photolysis).

69


3.2.2. Effect of different parameters on the photocatalytic degradation of LGSF dye for commercial TiO2 3.2.2.1. Effect of initial concentration of the dye Series of dye solutions were prepared with different concentrations ranging from 25 to 50 ppm in 100 mL with the addition of 0.175g of commercial TiO2 at 288.15 K and initial pH = 7.3. The results, listed in Table A-3 and depicted in Figure 3-23, show a pseudo first order reaction according to Langmuir Hinshelwood relationship. The results show that the apparent rate constant of the reaction decreases with the increase of the initial concentration of the dye. Besides, the relationship between (PDE) and time was listed in Table A-5 and plotted in Figure 3-24. 3.0 2.5

ln (C0/Ct)

2.0 30ppm 25ppm 40ppm 50ppm

1.5 1.0 0.5 0.0

0

10

20

30

t/min

40

50

60

70

Figure 3-23: The change of ln(C0/Ct) with irradiation time for different initial concentrations of LGSF dye solution using commercial TiO2.

70


100 90 80

PDE (%)

70 60

30ppm 25ppm 40ppm 50ppm

50 40 30 20 10 0

0

10

20

30

t/min

40

50

60

70

Figure 3-24: The relationship between PDE and irradiation time for different initial concentrations of LGSF dye solution using commercial TiO2.

3.2.2.2. Effect of the mass of commercial TiO2 A series of experiments were conducted using different masses of the commercial TiO2 added to the dye solution (30ppm in 100 mL with initial pH = 7.3 at 288.15 K). It was found, from these experiments, that the 0.3g of commercial TiO2/100 mL of LGSF dye gives the optimum photocatalytic activity. The results are listed in Table A-6 and plotted in Figure 3-25. Besides, (PDE) was calculated and the results were listed in Table A-8 and plotted in Figure 3-26.

71

Chapter Three: Results 3.0 2.5

ln(C0/Ct)

2.0

0.3 g 0.25 g

1.5

0.4 g 0.1 g

1.0

0.2 g

0.5 0.0

0.05 g 0

10

20

30

40

50

60

70

t/min

Figure 3-25: The change of ln(C0/Ct) with irradiation time for different masses of commercial TiO2. 100 90 80

PDE (%)

70 60 0.3 g 0.25 g 0.4 g 0.2 g 0.1 g 0.05 g

50 40 30 20 10 0

0

10

20

30

40

50

60

70

t/min

Figure 3-26: The relationship between PDE and irradiation time for different masses of commercial TiO2.

72

Chapter Three: Results 3.2.2.3. Effect of the initial pH of the solution The initial pH of the dye solution plays an essential role in the production of hydroxyl radicals, hence, a series of experiments were conducted at different initial pHs of the dye solution (30 ppm in 100 mL at 288.15 K) with the addition of 0.3 g of commercial TiO2 as a photocatalyst. The results, expressed in Table A-9 and plotted in Figure 3-27, show an increase in apparent rate constant of the reaction with the increase of the initial pH of the solution reaching a maximum level at pH = 7.3. The apparent rate constant decreases after this pH, so the optimum pH is found to be 7.3. Also, (PDE) was calculated and the results were listed in Table A-11 and plotted in Figure 3-28. 4 3.5 3

pH = 7.3

ln(C0/Ct)

2.5

pH = 6 pH = 8

2

pH = 5

1.5

pH = 9

1

pH = 4 pH = 3

0.5 0

0

10

20

30

40

t/min

Figure 3-27: The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of commercial TiO2.

73

Chapter Three: Results 100 pH = 7.3 pH = 6 pH = 8 pH = 5 pH = 9 pH = 4 pH = 3

PDE (%)

80 60 40 20 0

0

10

20

t/min

30

40

50

Figure 3-28: The relationship between PDE and irradiation time for different initial pH of LGSF dye solution using commercial TiO2.

3.2.2.4. Effect of the temperature To investigate the effect of temperature on photocatalytic degradation reaction of LGSF dye series of experiments were conducted under the range 283.15-308.15 K and using 30ppm in 100 mL of dye solution and 0.3g of the commercial TiO2 and initial pH equal to 7.3. The results, listed in Table A-12 and plotted in Figure 3-29, showed that the degradation rate increases with the increase of the temperature but for a certain limit. PDE was calculated also and the results were listed in Table A-17 and plotted in Figure 3-30. Activation energy of this reaction was calculated employing Arrhenius relationship (see Table A-14 and Figure 3-31), and it was found to be equal to 49.208 kJ mol–1. Thermodynamic parameters were obtained employing Eyring plot. The results were listed in Table A-15 and plotted in Figure 3-32.

74

Chapter Three: Results 3.5

ln(C0/Ct)

3 2.5

T = 293.15 K

2

T = 298.15 K

1.5

T = 288.15 K T = 303.15 K

1

T = 308.15 K

0.5 0

T = 283.15 K 0

20

40

60

t/min

Figure 3-29: The change of ln(C0/Ct) with irradiation time for different temperatures in the presence of commercial TiO2.

100 90 80

PDE (%)

70

T = 293.15

60

T = 298.15

50

T = 283.15

40

T = 288.15

30

T = 308.15

20

T = 303.15

10 0

0

10

20

30 t/min

40

50

60

70

Figure 3-30: The relationship between PDE and irradiation time for different temperatures using commercial TiO2.

75


0.0 -0.5 ln (k) min-1

-1.0 -1.5 -2.0 -2.5 -3.0 -3.5

3.3

3.35

3.4

3.45

3.5

3.55

(103/T)/K

Figure 3-31: Arrhenius relationship with commercial bulk TiO2.

-7.6

ln (k/T) (min-1K-1)

-7.8 -8.0 -8.2 -8.4 -8.6 -8.8 -9.0

3.3

3.35

3.4

3.45

3.5

(103/T)/K

Figure 3-32: Eyring plot of ln(k/T) vs. 1/T for commercial bulk TiO2.

76

3.55


3.2.3. Effect of different parameters on the photocatalytic degradation of LGSF dye for commercial nano TiO2 3.2.3.1. Effect of initial concentration of the dye A series of dye solutions was prepared with different concentrations ranging from 15 to 25 ppm in 100 mL with the addition of 0.175g of commercial nano TiO2 at 288.15 K and initial pH = 7.3. The results, listed in Table B-1 and depicted in Figure 3-33, show a pseudo first order reaction according to Langmuir Hinshelwood relationship. The results show that the apparent rate constant of the reaction decreases with the increase of the initial concentration of the dye. Besides, (PDE) was calculated and the results were listed in Table B-3 and plotted in Figure 3-34. 3.0 2.5

ln (C0/Ct)

2.0 1.5 1.0

20 ppm

0.5

25 ppm

0.0

15 ppm

0

10

20

30

40

50

60

70

t/min

Figure 3-33: The change of ln(C0/Ct) with irradiation time for different initial concentrations of LGSF dye solution using commercial nano TiO2.

77

Chapter Three: Results 100 90 80

PDE (%)

70 60 50 40

20 ppm

30

15 ppm

20

25 ppm

10 0

0

10

20

30

40

50

60

70

t/min

Figure 3-34: The relationship between PDE and irradiation time for different initial concentrations of LGSF dye solution using commercial nano TiO2. 3.2.3.2. Effect of the mass of commercial nano TiO2 A series of experiments were conducted using different masses of the commercial nano TiO2 added to the dye solution (20ppm in 100 mL with initial pH = 7.3 at 288.15 K). It was found, from these experiments, that the 0.2g of commercial nano TiO2/100 mL of LGSF dye gives the optimum photocatalytic activity. The results are listed in Table B-4 and plotted in Figure 3-35. Besides, (PDE) was calculated and the results were listed in Table B-6 and plotted in Figure 3-36.

78

Chapter Three: Results 4.5 4 3.5 0.2 g

3

0.15 g

ln(C0/Ct)

2.5

0.25 g

2

0.3 g

1.5

0.4 g

1

0.1 g 0.05 g

0.5 0

0

10

20

30

40

50

60

t/min

Figure 3-35: Variation of ln(C0/Ct) with irradiation time for different masses of

PDE (%)

commercial nano TiO2. 100 90 80 70 60 50 40 30 20 10 0

0.2 g 0.15 g 0.25 g 0.3 g 0.4 g 0.1 g 0

10

20

30

40

50

60

70

t/min

Figure 3-36: The relationship between PDE and irradiation time for different masses of commercial nano TiO2.

79

Chapter Three: Results 3.2.3.3. Effect of the initial pH of the solution The initial pH of the dye solution plays an essential role in the production of hydroxyl radicals, hence, a series of experiments were conducted at different initial pHs of the dye solution (20 ppm in 100 mL at 288.15 K) with the addition of 0.2 g of commercial TiO2 as a photocatalyst. The results, expressed in Table B-7 and plotted in Figure 3-37, show an increase in apparent rate constant of the reaction with the increase of the initial pH of the solution reaching a maximum level at pH = 7.3. The apparent rate constant decreases after this pH, so the optimum pH is found to be 7.3. Also, (PDE) was calculated and the results were listed in Table B-9 and plotted in Figure 3-38. 4 3.5

ln (C0/Ct)

3

pH = 7.3

2.5

pH = 6

2

pH = 5

1.5

pH = 8 pH = 9

1

pH = 4

0.5 0

pH = 3 0

10

20

30

40

50

t/min

Figure 3-37: The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of commercial nano TiO2.

80

Chapter Three: Results 100 90

PDE (%)

80 70

pH = 7.3

60

pH = 6

50

pH = 5

40

pH = 8

30

pH = 9

20

pH = 4

10 0

pH = 3 0

10

20

30 t/min

40

50

60

Figure 3-38: The relationship between PDE and irradiation time for different initial pH of LGSF dye solution using commercial nano TiO2. 3.2.3.4. Effect of the temperature The effect of temperature on the photocatalytic degradation reaction was investigated through a series of experiments conducted under the range 283.15308.15 K using 20ppm in 100 mL of dye solution and 0.2g of the commercial nano TiO2 and initial pH equal to 7.3. The results, listed in Table B-10 and plotted in Figure 3-39, showed that the degradation rate increases with the increase of the temperature but for a certain limit. PDE was calculated also and the results were listed in Table B-15 and plotted in Figure 3-40. The activation energy of this reaction was calculated employing Arrhenius relationship (see Table B-12 and Figure 3-41), and it was found to be equal to 31.912 kJ mol–1.

81

Chapter Three: Results Eyring plot was employed to obtain the thermodynamic parameters. The results are listed in Table B-13 and plotted in Figure 3-42 . 6 5 T = 283.15 K

ln (C0/Ct)

4

T = 288.15 K

3

T = 293.15 K

2

T = 298.15 K

1

T = 303.15 K

0

T = 308.15 K 0

20

40

60

80

t/min

Figure 3-39: The change of ln(C0/Ct) with irradiation time for different temperatures in the presence of commercial nano TiO2. 100 90 80 PDE (%)

70 60 T = 283.15 K T = 288.15 K T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K

50 40 30 20 10 0

0

10

20

30

40

50

60

70

t/min

Figure 3-40: The relationship between PDE and irradiation time for different temperatures using commercial nano TiO2.

82


0.0 -0.5 -1.0 ln(k)

-1.5 -2.0 -2.5 -3.0 -3.5 -4.0

3.25

3.3

3.35

3.4

3.45

3.5

3.55

(103/T)/K

Figure 3-41: Arrhenius relationship with commercial nano TiO2.

-8.0 -8.2

ln(k/T)

-8.4 -8.6 -8.8 -9.0 -9.2 -9.4

3.25

3.3

3.35

3.4 (103/T)/K

Figure 3-42: Eyring plot of ln(k/T) vs. 1/T.

83

3.45

3.5

3.55


3.3. Effect of calcination temperature on prepared TiO2 The effect of calcination temperature on prepared TiO2 for each mixing ratio of TiCl4 and ethanol was investigated under two different calcination temperatures 600 and 800 °C, and under experimental conditions of light intensity equals to 1.409×10–7 Einstein s–1, initial concentration of the dye solution equals 20 ppm, catalyst dosage equals to 0.2 g, initial pH equals to 7.3 and temperature equals to 298.15 K. The results of each calcination temperature are listed in Tables C-1 and plotted in Figure 3-43 and Figure 3-44 respectively. PDE was also calculated and the results were listed in Table C-3 and plotted in Figure 3-45 and Figure 3-46 respectively. It is found that the most efficient catalyst was the one that was prepared from the mixing ratio [V/V (TiCl4:EtOH) = 3:10] and calcinated at 600 °C. Figure 3-47 illustrates the variation of the apparent rate constant with the variation of the mixing ratio at 600 and 800° calcination temperatures. 3.5

5

3:10 1:10 1:4

3

2.5

ln (C0/Ct)

ln (C0/Ct)

4

2 1 0

1:4 1:10 3:10

3 2 1.5 1 0.5

0

0

5 10 15 20 25 30 35 40 45

0

10

t/min

20

30

40

50

60

70

t/min

Figure 3-43: Variation of ln(C0/Ct) with

Figure 3-44: Variation of ln(C0/Ct) with

irradiation time for different mixing

irradiation time for different mixing

ratios of the prepared TiO2 calcinated

ratios of the prepared TiO2 calcinated

under 600 °C.

under 800 °C.

84

100

100

80

80 PDE (%)

PDE (%)


60 3:10

40

1:10

20 0

60 40 20

1:4 0

1:4 1:10 3:10

0

5 10 15 20 25 30 35 40 45 t/min

0

10 20 30 40 50 60 70 t/min

Figure 3-46: The relationship between

PDE and irradiation time for different

PDE and irradiation time for different

mixing ratios of the prepared TiO2

mixing ratios of the prepared TiO2

calcinated under 600 °C.

calcinated under 800 °C.

k/min-1


0.14 0.12 0.1 0.08 0.06 0.04 0.02 0

600 °C 800 °C

1:4

3:10

1:10

V/V ratio

Figure 3-47: The variation of the apparent rate constant with the variation of the mixing ratio at 600 and 800° calcination temperatures

85


3.4. Effect

of

different

parameters

on

the

photocatalytic

degradation of LGSF dye with prepared TiO2 and calcination at 600 °C 3.4.1. Effect of initial dye concentration The effect of initial dye concentration was investigated under identical conditions in comparison with the experiments conducted for the commercial nano TiO2. The results were listed in Tabel-C-4 and plotted in Figure 3-48. In addition, PDE was calculated and the results were listed in Table C-6 and plotted in Figure 3-49. 4 3.5 3

ln(C0/Ct)

2.5 2 1.5

15 ppm 20 ppm 25 pmm

1 0.5 0

0

10

20

30

40

t/min

Figure 3-48: The change of ln(C0/Ct) with irradiation time for different initial concentrations of LGSF dye solution using prepared TiO2.

86

Chapter Three: Results 100 90 80 PDE (%)

70 60 50 40

15 ppm 20 ppm 25 pmm

30 20 10 0

0

5

10

15

20 t/min

25

30

35

40

Figure 3-49:The relationship between PDE and irradiation time for different initial concentrations of LGSF dye using prepared TiO2.

3.4.2. Effect of the mass of prepared TiO2 The effect of the mass of the prepared TiO2 was investigated under identical conditions in comparison with the experiments conducted for the commercial nano TiO2. The results were listed in Table C-7 and plotted in Figure 3-48. In addition, PDE was calculated and the results were listed in Table C-9 and plotted in Figure 3-49. It was found, from these experiments, that the 0.2g of prepared TiO2/100 mL of LGSF dye gives the optimum photocatalytic activity.

87

Chapter Three: Results 8 7 6

0.2 g

ln(C0/Ct)

5

0.25 g

4

0.15 g

3

0.3 g 0.4 g

2

0.1 g

1 0

0.05 g 0

10

20

30

40

50

60

t/min

Figure 3-50: The variation of ln(C0/Ct) with irradiation time for different

PDE (%)

masses of the prepared TiO2.

100 90 80 70 60 50 40 30 20 10 0

0.2 g 0.25 g 0.15 g 0.3 g 0.4 g 0.1 g 0.05 g 0

10

20

30

40 t/min

50

60

70

80

Figure 3-51: The relationship between PDE and irradiation time for different masses of prepared TiO2.

88


3.4.3. Effect of initial pH of the solution The effect of initial pH of the solution was investigated under identical conditions in comparison with the experiments conducted for the commercial nano TiO2. The results were listed in Table C-10 and plotted in Figure 3-48. In addition, PDE was calculated and the results were listed in Table C-12 and plotted in Figure 3-49. The results show an increase in apparent rate constant of the reaction with the increase of the initial pH of the solution reaching a maximum level at pH = 7.3. The apparent rate constant decreases after this pH, so the optimum pH is found to be 7.3. 7 6 pH = 7.3

5

pH = 6

ln (C0/Ct)

4

pH = 5

3

pH = 4 pH = 3

2

pH = 8

1 0

pH = 9 0

10

20

30

40

50

t/min

Figure 3-52: The variation of ln (C0/Ct) with irradiation time for different initial pHs of the dye solution.

89


100 90 80

pH = 7.3

PDE (%)

70

pH = 6

60 50

pH = 5

40

pH = 4

30

pH = 3

20

pH = 8

10

pH = 9

0

0

10

20

30 t/min

40

50

60

Figure 3-53: The relationship between PDE and irradiation time for different initial pHs of the dye solution.

3.4.4. Effect of temperature on dye removal The effect of temperature on the photocatalytic degradation reaction of was studied by conducting a series of experiments under temperature range 283.15308.15 K and using 20ppm in 100 mL of dye solution and 0.2g of the prepared TiO2 and initial pH equal to 7.3. The results, listed in Table C-13 and plotted in Figure 3-54, showed that the degradation rate increases with the increase of the temperature but for a certain limit. PDE was calculated also and the results were listed in Table C-18 and plotted in Figure 3-55. The activation energy of this reaction was calculated employing Arrhenius relationship, and from the results, which are listed in Table C-15 and plotted in Figure 3-56, the activation energy was found to be equal to 65.977 kJ mol–1. The

90

Chapter Three: Results thermodynamic parameters were obtained using Eyring plot, and the results were

ln(C0/Ct)

listed in Table C-16 and plotted in Figure 3-57. 10 9 8 7 6 5 4 3 2 1 0

T = 303.15 T = 298.15 T = 308.15 T = 293.15 T = 288.15 T = 283.15 0

20

40

60

80

t/min

Figure 3-54: The change of ln(C0/Ct) with irradiation time at different

PDE (%)

temperatures using prepared TiO2 as a catalyst. 100 90 80 70 60 50 40 30 20 10 0

T = 303.15 T = 298.15 T = 308.15 T = 293.15 T = 288.15 T = 283.15 0

10

20

30

40

50

60

70

t/min

Figure 3-55: The relationship between PDE and irradiation time at different temperatures using prepared TiO2 as a catalyst

91


0.0 -0.5 ln(k) (min-1)

-1.0 -1.5 -2.0 -2.5 -3.0 -3.5 -4.0

3.25

3.3

3.35

3.4

3.45

3.5

3.55

(103/T)/K

ln (k/T) (min-1 K-1)

Figure 3-56: Arrhenius relationship with prepared TiO2.

-7.6 -7.8 -8.0 -8.2 -8.4 -8.6 -8.8 -9.0 -9.2 -9.4

3.25

3.3

3.35

3.4 (103/T)/K

Figure 3-57: Eyring plot with prepared TiO2.

92

3.45

3.5

3.55

Chapter Four

Discussion

Chapter Four: Discussion

4.1. Characterization of the catalysts The three types of TiO2 used in this work were characterized employing XRD, AFM and Fluorescence spectroscopy techniques.

4.1.1. X-Ray Diffraction patterns (XRD) The XRD patterns of the three types of TiO2 used in this work are shown in Figures Figure 3-1 to Figure 3-8 respectively. The characteristic diffraction peaks of the anatase phase of TiO2 at 2θ = 25.25° (101) and 48.0° (200)[96] appear in the patterns, which means that anatase is the primary phase in the catalysts used for the degradation of LGSF dye. Clearly, there are no peaks of either rutile or brookite phase in the patterns of the samples calcinated at 600°C, and this is an evidence of the high purity of the prepared samples. The XRD patterns showed the transformation from anatase to rutile phase occurred in the samples calcinated at 800°C, since XRD patterns exhibited strong diffraction peaks at 2θ = 27°, 36° and 55° indicating TiO2 in rutile phase. Additionally, the XRD patterns revealed that the optimum prepared TiO2 nanoparticle for the degradation of LGSF is pure anatase with no rutile or brookite phase exist.

4.1.2. Atomic Force Microscopy (AFM) AFM images showed that particles of the prepared titanium dioxide have spherical shape.

4.1.3. Fluorescence spectroscopy The results obtained from fluorescence spectroscopy showed that the bands of the photocatalysts used in this study suffered from blue shift which leads to an increase in band gap energy (3.51 – 3.56 eV), however, these values are slightly larger than the reported value (3.3 – 3.4 eV)[97]. 93


4.2. Preliminary Experiments Series of experiments were conducted and the results are shown in Figure 4-1. 8.0

Photocatalysis with prepared TiO2

7.0

Photocatalysis with commercial bulk TiO2

ln (C0/Ct)

6.0 5.0

Photocatalysis with commercial nano TiO2

4.0

Photolysis

3.0

Dark reaction

2.0 1.0 0.0

0

5 10 15 20 25 30 35 40 45 50 55 60 65 t/min

Figure 4-1: Preliminary experiments with commercial bulk, commercial nano and prepared TiO2. In the absence of UV-light and under O2, no reaction took place since there was no (e– - h+) pairs generated, moreover, no reaction took place in the absence of the catalysts. (dark reaction)

TiO2 comm. or comm. nano or prepared + Dye + O2 �⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯� N.R hυ (photolysis)

Dye + O2 �⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯� N.R

4-1 4-2

On the other hand, the reaction proceeded in the presence of both catalyst and UV-light, as follows: hυ

TiO2 comm, comm nano, prepared nano + Dye + O2 �⎯⎯⎯⎯⎯⎯� P

4-3

This shows clearly that the presence of light, electron scavenger (O2) and photocatalyst (TiO2) are essential for effectual photocatalytic degradation of LGSF dye.

94


4.3. Effect of different parameters on photocatalytic degradation of LGSF dye using commercial, commercial and prepared nano TiO2 This work intends to determine the effect of different parameters on photocatalytic degradation of LGSF dye using different types of TiO2 as photocatalyst.

4.3.1. Effect of initial dye concentration Generally, wastewater produced from industries or laboratories contain different amount of pollutants, hence the investigation of the effect of initial dye concentration is an important requirement for the successful application of the photocatalytic degradation system. The effect of initial LGSF concentration on the rate of the photocatalytic degradation reaction in the presence of commercial, commercial nano and prepared TiO2 was investigated in the range of (15-50 ppm) and under fixed conditions. The results, expressed in Figure 4-2, show that the efficiency of the photocatalytic degradation reaction increased with the increase of the initial concentration of LGSF dye and reached its maximum value at 20 ppm for commercial nano and prepared TiO2, but 30 ppm for the commercial bulk TiO2. This behavior can be rationalized on the basis that as the initial dye concentration increases, larger amounts of dye molecules will be adsorbed on the surface of TiO2 and this decrease the number of photons that are capable to reach the surface of the catalyst, hence less OH radicals are produced, declining the degradation process[25].

95

Chapter Four: Discussion 0.12 0.1

prepared TiO2 commercial bulk TiO2

k/min-1

0.08

commercial nano TiO2

0.06 0.04 0.02 0

0

10

20 30 40 initial concentration of LGSF dye/ppm

50

60

Figure 4-2: The relationship between the apparent rate constant and the initial concentration of LGSF dye. PDE results showed that the efficiency of the optimal initial concentration of LGSF dye in the presence of commercial bulk and commercial nano and prepared TiO2 reached 82.622% and 57.377% and 91.703% after 25 min respectively, which indicates that the prepared catalyst is economically more beneficial than the other two types.

4.3.2. Effect of the mass of the catalyst The amount of the catalyst used in the photocatalytic degradation reaction is one of the influential factor on the efficiency of this reaction, so series of experiments were conducted to determine the optimum amount of the catalyst for the three types of TiO2 used in this study. The results, plotted in Figure 4-3, showed that the photodegradation efficiency increased with the increase of the amount of the catalyst reaching a maximum

96

Chapter Four: Discussion value at 0.3 g /100 mL for the commercial TiO2 and 0.2 g / 100 mL for commercial nano and prepared TiO2. 0.14

prepared TiO2

0.12

commercial TiO2 commercial nano TiO2

k /min-1

0.1 0.08 0.06 0.04 0.02 0

0

0.1

0.2

0.3

0.4

0.5

Dosage of the catalyst/g

Figure 4-3: The relationship between the apparent rate constant and the dosage of the catalyst. This behavior can be rationalized by three possibilities[98]: 1- The additional amount of the photocatalyst would no more enhance the degradation efficiency when all dye molecules are adsorbed on photocatalyst surface. 2- As the number of particles in the solution increases, the particle-particle interaction becomes more significant, and this may lead to an increase in the rate of the deactivation of activated molecules through collision with ground state TiO2 particles. 3- The excess in the photocatalyst particles may lead to an increase in the opacity of the suspension which may cause retardation to degradation rate due to the blocking of light penetration by the large amounts of the photocatalyst.

97

Chapter Four: Discussion PDE results showed that the efficiency of using 0.2 g of commercial bulk and commercial nano and prepared TiO2 reached 84.685% and 72.806% and 94.730% after 25 min respectively, which indicates that the prepared catalyst is economically more beneficial than the other two types.

4.3.3. Effect of pH on dye removal The pH of the solution is one of the most important parameters for the photocatalytic degradation of dyes, since it may affect dye adsorption onto the surface of the semiconductor as the charge of the catalyst surface is strongly influenced by pH[99]. The examined range of pH was (3-9), and in all the experiments, the adjustment of pH was carried out by the addition of appropriate amount of NaOH or HCl aqueous solutions. The results of the experiments are plotted in Figure 4-4. These results show clearly that the efficiency of the process depends strongly on the pH of the solution, as the value of the apparent rate constant increase with the increase of pH and reach its maximum at pH = 7.3 for all of the studied catalysts. The efficiency decreased before and after pH = 7.3, and this depends on the point of zero charge (pHpzc), which is a concept related to adsorption phenomenon, and it is a description of the condition when the electrical charge density on a surface is equal to zero[100]. Under acidic or basic conditions the surface of TiO2 is protonated or deprotonated respectively, according to the following equations[25, 98]: TiOH(surface) + H+ �⎯⎯⎯� TiOH+2(surface)

TiOH(surface) + OH– �⎯⎯⎯� TiO–(surface) + H2 O

98

4-4 4-5

Chapter Four: Discussion prepared TiO2 commercial bulk TiO2 commercial nano TiO2

0.14 0.12

k/min-1

0.10 0.08 0.06 0.04 0.02 0.00

2

3

4

5 6 7 Initial pH of the dye solution

8

9

10

Figure 4-4: The relationship between the apparent rate constant and pH of dye solution. This means that TiO2 surface possesses positive charge (attracts anions) in acidic media, and negative charge (attracts cations) in basic media. The adsorption of the dye molecules onto the catalyst surface, which is a significant step for the photocatalytic oxidation to occur, is influenced by the change in pH value[101]. In basic medium, there exist a Coulombic repulsion between the negative charged surface of photocatalyst and OH– anions, so the formation of ●OH will decline and thus decrease the photooxidation. As the LGSF dye molecules are negatively charged in alkaline medium, it is expected that due to the Coulombic repulsion with TiO– groups on the surface of the photocatalyst, LGSF molecules will be adsorbed scarcely and the photooxidation declines[102, 103].

99

Chapter Four: Discussion PDE results showed that the efficiency of the degradation in pH = 7.3 for commercial bulk and commercial nano and prepared TiO2 reached 82.119% and 72.806% and 94.730% respectively after 25 minutes.

4.3.4. Effect of temperature on dye removal The effect of temperature on the degradation rate of LGSF dye was investigated by a series of experiments and the examined range was (283.15 - 308.15) plotted in Figure 4-5. 0.14 0.12

k/min-1

0.1 0.08 0.06 0.04

Prepared TiO2

0.02

Commercial bulk TiO2

0

Commercial nano TiO2 280

285

290

295

300

305

310

T/K

Figure 4-5: The relationship between the apparent rate constant and temperature of dye solution. Increasing the temperature of the reaction increases the rates of all reactions taking place in the system with the exception of electron-hole pairs photogeneration[104], but the increase in temperature will lower the solubility of oxygen in the solution leading to a decrease in the rate of electron withdrawal from the surface of TiO2[105]. In addition, increase of temperature will lead to an increase in the recombination rate of charge carriers and the desorption of the

100

Chapter Four: Discussion adsorbed reactants onto the surface of the TiO2, and this will lead to a decline in the activity of photocatalysis process[25, 104]. The investigation of activation energy values for the studied photocatalysts showed that the commercial bulk TiO2 has the largest activation energy (48.676 kJ mol–1) in comparison with prepared and commercial nano TiO2 (45.780 and 37.032 kJ mol–1 respectively). The obtained results are in agreement with the average particle sizes obtained by AFM, which show that the average particle size of the commercial bulk TiO2 is the largest among the used photocatalysts (44.0 nm), followed by the prepared TiO2 (29.3 nm) and the commercial nano TiO2 (18.0 nm). As the average particle size becomes smaller, the adsorption of the relatively large dye molecules become harder, hence, the stability of the dye molecule decreases leading to a consequent decrease in the activation energy. The results plotted in Figure 4-6 and Figure 4-7 and listed in Tables A-16 and B-14 and C-17, showed also that the photocatalytic degradation of LGSF is an endothermic process, and the low value of the entropy is attributed to decrease in randomness. In addition, the positive values of Gibb’s free energy ∆G are an evident that the photocatalytic degradation of LGSF dye is not spontaneous. PDE results showed that the efficiency of the degradation under 303.15 K for commercial bulk and commercial nano and prepared TiO2 reached 73.563% and 69.531% and 94.730% after 25 min respectively.

101


Commercial bulk TiO2

0.0

Prepared TiO2

-0.5

ln (k/T)

-1.5 -2.0 -2.5 -3.0

-7.0 -8.0 -9.0

-3.5 -4.0

Commercial bulk TiO2 Prepared TiO2 Commercial nano TiO2

-6.0

Commercial nano TiO2

-1.0 ln k

-5.0

-10.0 3.2

3.4

(103/T)/K

3.6

3.2

3.4

3.6

(103/T)/K

Figure 4-6: The relationship between lnk


and 1/T.

ln(k/T) and 1/T.

4.4. Suggested mechanism for dye removal The main source of OH radicals is the reaction between the holes and the surface-adsorbed water or hydroxyl ions. It is presumed that degradation of LGSF dye occurs mostly via attack by OH radicals on the N-ethyl groups of LGSF dyes under neutral and basic conditions, and through attack of OH radicals on the central carbon atom of LGSF to destruct the conjugated structure of the dye[106], as follows:

102

Chapter Four: Discussion O

O

N

S O O

O S O

O

O

N

S O O

O S O

O

O

N

S O O

O S O

-H2O O

N

S O

O

O

OH

O

H

N

S O

O

OH

O

N

S O

OH

-CH3CHO

H O

O

N

S O O

O S O

O

O

N

S O O

O S O

O

O

N

S O O

O S O

OH -H2O O O

N

S O

O H

O

O

N

S O

H

N

S O

O

H

OH

HO O

O

N

S O O

O O S O

H

N

O

O S O

H

N

O S O

OH -CH3CHO

OH O

O

O

N

S O O

S O

N H

S O O

H

O

O

O N H

S O O

O S O

O2 OH H

O O S O

N

O H

O S O O

NH

+

O O

O S O O

+ 2OH

N

O O

O S O

OH O

O N H

S O O

H2O

O S O

OH +

O O

CHO

S O

H

NH2

O O S O

N

a

H

O O S O

N

+ H 2O 2

OH O S O O

O S O O

O OH

O

+ 2OH further degradation

mineralization products

a + b

Scheme 4-1: Suggested mechanism for the degradation of LGSF dye. Adapted from reference[106].

103

Chapter Five

Conclusions, Recommendations And references

Chapter Five: Conclusions, Recommendations and References

5.1. Conclusions This study investigated the efficiency of commercial bulk, commercial nano and prepared nano TiO2 as a photocatalyst in the degradation of LGSF dye solution, and emerged conclusions can be summarized as follows: 1. The mean crystallite size and crystallite size for the commercial bulk, commercial nano and prepared TiO2 were calculated using Scherrer and modified Scherrer equation, and the results showed that the average crystallite size and average particle size of all photocatalysts were ranging from 27.0 to 48.1 nm and 66.72 to 432.93 nm according to XRD and AFM techniques respectively. 2. XRD patterns showed also that the transformation from anatase to rutile phase occurred in the samples calcinated at 800°C, while no rutile phase presents in the samples calcinated at 600°C. 3. AFM images showed that particles of the prepared titanium dioxide have spherical shape, and that the particles sizes of the prepared samples were found to be larger than the values of crystallite size and average crystallite size. 4. Fluorescence spectra of commercial bulk, commercial nano and prepared TiO2 showed that the band gap energy of the prepared TiO2 is in the range (3.51-3.56 eV) which is slightly above the reported value. 5. The best calcination temperature was found to be 600 °C. 6. The best mixing ratio was found to be [V/V (TiCl4:EtOH) = 3:10]. 7. The best initial concentration of the dye was found to be 30 ppm for commercial bulk TiO2, and 20ppm for commercial and prepared nano TiO2.

104

Chapter Five: Conclusions, Recommendations and References 8. The maximum catalyst dose was found to be (0.3 g / 100 mL) for commercial bulk TiO2, and (0.2 g / 100 mL) for commercial and prepared nano TiO2. These results substantiate the economic usefulness of the commercial nano and prepared TiO2. 9. The optimum pH of the solution was found to be 7.3, and the efficiency of the photocatalyst decreased below and above this pH. 10.The activation energy of the commercial bulk TiO2 was found to be largest amongst the the studied photocatalysts. The values of the activation energy are in good agreement with the results obtained from AFM technique. 11.Using Eyring equation thermodynamic parameters (∆H, ∆S) were caluculated, and the results showed that the reaction is endothermic with low randomness. ∆G was calculated via Gibb’s equation and the results showed that the reaction is non-spontaneous.

5.2. Recommendations 1. Loading several metals like Pt, Pd, Ni, Co, with different percentages on the surface of TiO2 may enhance the efficiency of TiO2 nanoparticles. 2. Synthesis of carbon nanotube-titanium dioxide nanoparticles would be interesting, as carbon nanotubes could anchor TiO2 nanoparticles and increase their efficiency. 3. Studying the effect of other parameters, like flow rate of O2 and type of the current gas, is very important. 4. It would be interesting to prepare TiO2 nanoparticles via other preparation methods and study the efficiency of the resultant nanoparticles in the degradation of LGSF dye.

105

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116

Appendix A Table A-1: The change of ln(C0/Ct) with time in the absence of radiation (Dark reaction). ln(C0/Ct) time (min) Commercial Commercial

0 5 10 15 20 25 30 35 40 45 50 55 60

TiO2

nano TiO2

0.000 0.000 0.653E-03 0.000 0.634E-03 1.961E-03 1.307E-03 1.307E-03 1.961E-03 1.307E-03 2.616E-03 2.616E-03 1.961E-03

0.000 0.678E-03 0.678E-03 1.358E-03 1.358E-03 2.037E-03 2.037E-03 2.717E-03 2.717E-03 2.717E-03 3.398E-03 3.398E-03 4.079E-03

Table A-2: The change of ln(C0/Ct) with irradiation time in the absence of the catalyst (photolysis). Irradiation time (min) 0 5 10 15 20 25 30

ln(C0/Ct) 0.000 0.006 0.011 0.017 0.023 0.031 0.038

Irradiation time (min) 35 40 45 50 55 60 -

117

ln(C0/Ct) 0.051 0.061 0.068 0.075 0.087 0.099 -

Appendix A Table A-3: The change of ln(C0/Ct) with irradiation time for different initial concentrations of LGSF dye solution using commercial TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

ln(C0/Ct) for different concentrations of dye solution 25ppm 30ppm 40ppm 50ppm 0.000 0.000 0.000 0.000 0.324 0.292 0.233 0.168 0.620 0.640 0.412 0.337 0.822 0.997 0.651 0.500 1.008 1.476 0.843 0.760 1.303 1.750 0.944 0.944 1.659 1.986 1.156 1.148 1.894 2.104 1.334 1.367 2.140 2.238 1.523 1.598 2.553 2.327 1.726 1.673 1.990 1.816 2.159 2.020 2.476 2.222 -

Table A-4: The relationship between the apparent rate constant and the initial concentration of LGSF dye solutions using commercial TiO2. LGSF conc. (ppm)

k (min–1)

25 30 40 50

0.0547 0.0594 0.0402 0.0374

118

Appendix A Table A-5: The relationship between PDE and irradiation time for different initial concentrations of LGSF dye solution using commercial TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

PDE for different concentrations of dye solution (%) 25ppm 30ppm 40ppm 50ppm 0.000 0.000 0.000 0.000 27.682 25.305 20.798 15.453 46.194 47.256 33.761 28.641 56.055 63.110 47.865 39.320 63.495 77.134 56.978 53.236 72.837 82.622 61.111 61.084 80.969 86.280 68.519 68.285 84.948 87.805 73.647 74.515 88.235 89.329 78.205 79.773 92.215 90.244 82.194 81.230 86.325 83.738 88.462 86.731 91.595 89.159

Table A-6: The change of ln(C0/Ct) with irradiation time for different masses of commercial TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

0.05g 0.000 0.105 0.213 0.304 0.423 0.525 0.627 0.724 0.827 0.938 1.050 1.263 1.342

ln(C0/Ct) for different masses 0.1g 0.2g 0.25g 0.3g 0.000 0.000 0.000 0.000 0.228 0.370 0.457 0.443 0.373 0.655 0.830 0.828 0.585 0.903 1.047 1.280 0.759 1.311 1.499 1.721 0.918 1.477 2.058 1.063 1.876 1.258 1.485 1.707 2.039 2.229 2.407 -

119

0.4g 0.000 0.462 0.760 1.067 1.399 1.818 -

Appendix A Table A-7: The relationship between the apparent rate constant and the mass of the commercial TiO2. Catalyst dosage (g)

k (min–1)

0.05 0.1 0.2 0.25 0.3 0.4

0.0246 0.0389 0.0604 0.0751 0.0855 0.0719

Table A-8: The relationship between PDE and irradiation time for different masses of commercial TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

PDE for different masses of the catalyst (%) 0.05g 0.1g 0.2g 0.25g 0.3g 0.4g 0.000 0.000 0.000 0.000 0.000 0.000 9.964 20.400 30.931 36.702 35.762 37.013 19.210 31.164 48.048 56.383 56.291 53.247 26.212 44.305 59.459 64.894 72.185 65.584 34.470 53.191 73.045 80.407 82.119 75.325 40.844 60.075 77.177 87.234 83.766 46.589 65.457 84.685 51.526 71.589 56.284 77.347 60.862 81.852 64.991 86.984 71.724 89.237 73.878 90.989 -

120

Appendix A Table A-9: The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of commercial TiO2. t (min)

ln(C0/Ct) at different pH 3

4

5

6

7.3

8

9

0

0.000

0.000

0.000

0.000

0.000

0.000

0.000

5

0.251

0.315

0.265

0.378

0.443

0.331

0.221

10

0.449

0.509

0.568

0.706

0.828

0.664

0.469

15

0.670

0.749

0.842

1.062

1.280

1.013

0.706

20

0.928

0.944

1.296

1.533

1.242

1.040

25

1.231

1.274

1.667

2.129

1.721 -

1.861

1.332

30

1.353

1.613

1.896

-

-

-

1.733

35

1.504

-

-

-

-

-

2.087

40

1.784

-

-

-

-

-

-

Table A-10: The relationship between the apparent rate constant and the initial pH of LGSF dye solution using commercial TiO2. pH

k (min–1)

3 4 5 6 7.3 8 9

0.0450 0.0514 0.0633 0.0791 0.0855 0.0690 0.0559

121

Appendix A Table A-11: The relationship between PDE and irradiation time for different initial pH of LGSF dye solution using commercial TiO2. PDE for different pHs (%)

t (min)

3

4

5

6

7.3

8

9

0

0.000

0.000

0.000

0.000

0.000

0.000

0.000

5

22.222 27.010 23.245 31.498 35.762 28.148 19.789

10

36.176 39.871 43.341 50.661 56.291 48.519 37.467

15

48.837 52.733 56.901 65.419 72.185 63.704 50.660

20

60.465 61.093 72.639 78.414 82.119 71.111 64.644

25

70.801 72.026 81.114 88.106

30

74.160 80.064 84.988

-

-

-

82.322

35

77.778

-

-

-

-

-

87.599

40

83.204

-

-

-

-

-

-

-

84.444 73.615

Table A-12: The change of ln(C0/Ct) with irradiation time for different temperatures in the presence of commercial TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

ln(C0/Ct) for different temperatures (K) 283.15 0.000 0.278 0.457 0.646 0.790 1.008 1.180 1.424 1.638 1.837 2.102 2.360 2.530

288.15 0.000 0.180 0.396 0.636 0.953 1.429 1.804 2.131 2.773 2.878 -

293.15 0.000 0.443 0.828 1.280 1.721 -

122

298.15

303.15

308.15

0.000 0.266 0.833 1.121 1.526 1.689 -

0.000 0.261 0.626 0.802 1.247 1.330 -

0.000 0.188 0.419 0.617 0.930 1.112 1.286 -

Appendix A Table A-13: The relationship between the apparent rate constant and the temperature of LGSF dye solution using commercial TiO2. T (K)

k (min–1)

283.15 288.15 293.15 298.15 303.15 308.15

0.0416 0.0615 0.0848 0.1202 0.0943 0.0713

Table A-14: The variation of 1/T with lnk. (103/T) (K)

lnk (min–1)

3.53 3.47 3.41 3.35 3.29 3.24

-3.180 -2.789 -2.459 -2.119 -2.361 -2.641

Table A-15: The variation of 1/T with ln(k/T). (103/T) (K)

ln(k/T) (min–1 K–1)

3.53 3.47 3.41 3.35 3.29 3.24

-8.826 -8.452 -8.140 -7.816 -8.076 -8.371

123

Appendix A Table A-16: Thermodynamic parameters of the decolorization of LGSF dye using commercial TiO2. Ea (kJ mol–1)

∆H‡ (kJ mol–1)

48.676

29.603

∆S‡ (kJ mol–1 K–1) ∆G‡288.15 (kJ mol–1) -0.107

78.847

Table A-17: The relationship between PDE and irradiation time for different temperatures using commercial TiO2. t (min)

PDE for different temperatures (K) 283.15

288.15

293.15

298.15

303.15

308.15

0

0.000

0.000

0.000

0.000

0.000

0.000

5

24.259

16.458

35.762

23.370

22.989

17.105

10

36.667

32.708

56.291

56.522

46.552

34.211

15

47.593

47.083

72.185

67.391

55.172

46.053

20

54.630

61.458

82.119

78.261

71.264

60.526

25

63.519

76.042

-

81.522

73.563

67.105

30

69.259

83.542

-

-

-

72.368

35

75.926

88.125

-

-

-

-

40

80.556

93.750

-

-

-

-

45

84.074

94.375

-

-

-

-

50

87.778

-

-

-

-

-

55

90.556

-

-

-

-

-

60

92.037

-

-

-

-

-

124

Appendix (B)

Appendix B Table B-1: The change of ln(C0/Ct) with irradiation time for different concentrations of LGSF dye solution using commercial nano TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

ln(C0/Ct) for different concentrations of dye solution (ppm) 15 20 25 0.000 0.000 0.000 0.086 0.164 0.055 0.228 0.321 0.139 0.415 0.482 0.237 0.509 0.650 0.342 0.748 0.853 0.426 0.889 1.090 0.604 1.136 1.249 0.755 1.317 1.767 0.949 1.667 1.825 1.089 1.778 2.096 1.382 2.057 2.153 1.703 2.268 2.607 1.892

Table B-2: The relationship between the apparent rate constant and the initial concentration of LGSF dye solutions using commercial nano TiO2. LGSF conc. (ppm)

k (min–1)

15 20 25

0.0351 0.0403 0.0212

125

Appendix B Table B-3: The relationship between PDE and irradiation time for different initial concentrations of LGSF dye solution using commercial nano TiO2. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

PDE for different concentrations of dye solution (%) 15ppm 20ppm 25ppm 0.000 0.000 0.000 8.219 15.164 5.318 20.396 27.459 12.990 33.942 38.251 21.099 39.878 47.814 28.945 52.664 57.377 34.699 58.904 66.393 45.336 67.884 71.311 53.008 73.212 82.923 61.290 81.126 83.880 66.347 83.105 87.705 74.891 87.215 88.388 81.779 89.650 92.623 84.917

Table B-4: The change of ln(C0/Ct) with irradiation time for different masses of commercial nano TiO2. t (min)

0 5 10 15 20 25 30 35 40 45 50 55 60

0.05 0.000 0.027 0.081 0.095 0.157 0.215 0.310 0.368 0.478 0.551 0.643 0.822 0.955

ln(C0/Ct) for different masses (g) 0.1 0.15 0.2 0.25 0.3 0.000 0.000 0.000 0.000 0.000 0.094 0.195 0.184 0.175 0.126 0.196 0.450 0.402 0.411 0.286 0.313 0.730 0.807 0.600 0.462 0.491 0.973 0.949 0.841 0.649 0.655 1.441 1.302 1.126 0.898 0.869 1.705 1.654 1.388 1.134 1.165 1.859 2.033 1.808 1.357 1.464 2.712 2.694 1.960 1.534 1.595 2.958 3.177 2.155 1.899 1.726 3.597 3.548 2.774 2.148 1.978 3.385 2.361 2.351 3.845 2.497

126

0.4 0.000 0.121 0.298 0.473 0.635 0.794 0.989 1.225 1.375 1.648 1.952 2.393 2.807

Appendix B Table B-5: The relationship between the apparent rate constant and the mass of the commercial TiO2. Catalyst dosage (g)

k (min–1)

0.05 0.1 0.15 0.2 0.25 0.3 0.4

0.0130 0.0370 0.0633 0.0551 0.0407 0.0395 0.0130

Table B-6: The relationship between PDE and irradiation time for different masses of commercial nano TiO2. t (min)

0.05

0 5 10 15 20 25 30 35 40 45 50 55 60

0.000 2.634 7.741 9.018 14.525 19.314 26.656 30.806 37.989 42.378 47.406 56.026 61.532

PDE for different masses (%) 0.1 0.15 0.2 0.25 0.3 0.000 8.957 17.800 26.871 38.776 48.073 58.050 68.821 76.871 79.705 82.200 86.168 90.476

0.000 17.749 36.219 51.804 62.193 76.335 81.818 84.416 93.362 94.805 97.258 -

0.000 16.835 33.094 55.396 61.295 72.806 80.863 86.906 93.237 95.827 97.122 -

127

0.000 16.043 33.690 45.098 56.863 67.558 75.045 83.601 85.918 88.414 93.761 96.613 97.861

0.000 11.826 24.850 36.976 47.754 59.281 67.814 74.251 78.443 85.030 88.323 90.569 91.766

0.4 0.000 11.419 25.775 37.684 46.982 54.812 62.806 70.636 74.715 80.750 85.808 90.865 93.964

Appendix B Table B-7: The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of commercial nano TiO2. t (min)

ln(C0/Ct) at different pH 3

4

5

6

7.3

8

9

0

0.000

0.000

0.000

0.000

0.000

0.000

0.000

5

0.156

0.249

0.143

0.217

0.184

0.193

0.189

10

0.357

0.413

0.461

0.478

0.402

0.367

0.445

15

0.486

0.595

0.650

0.847

0.807

0.688

0.651

20

0.555

0.876

0.925

1.303

0.949

0.999

0.885

25

0.708

1.053

1.252

1.549

1.302

1.280

1.212

30

0.993

1.511

1.777

-

1.654

1.782

1.470

35

1.144

1.997

-

2.033

40

1.424

1.613 -

2.508

-

2.694

2.145 -

1.722 -

Table B-8: The relationship between the apparent rate constant and the initial pH of LGSF dye solution using commercial nano TiO2. pH

k (min–1)

3 4 5 6 7.3 8 9

0.0328 0.0457 0.0566 0.0611 0.0642 0.0558 0.0479

128

Appendix B Table B-9: The relationship between PDE and irradiation time for different initial pH of LGSF dye solution using commercial nano TiO2. PDE for different pH (%)

t (min)

3

4

5

6

7.3

8

9

0 5 10 15 20 25 30 35 40

0.000 14.444 30.000 38.519 42.593 50.741 62.963 68.148 75.926

0.000 22.064 33.808 44.840 58.363 65.125 77.936 80.071 -

0.000 13.361 36.952 47.808 60.334 71.399 83.090 86.430 91.858

0.000 19.512 37.979 57.143 72.822 78.746 84.321 -

0.000 16.835 33.094 55.396 61.295 72.806 80.863 86.906 93.237

0.000 17.561 30.732 49.756 63.171 72.195 83.171 88.293 -

0.000 17.248 35.934 47.844 58.727 70.226 77.002 82.136 -

Table B-10: The change of ln(C0/Ct) with irradiation time for different temperatures in the presence of commercial nano TiO2. t (min)

ln(C0/Ct) for different temperatures (K) 283.15

288.15

293.15

298.15

303.15

308.15

0

0.000

0.000

0.000

0.000

0.000

0.000

5

0.056

0.107

0.207

0.207

0.170

0.209

10

0.122

0.283

0.461

0.460

0.458

0.404

15

0.168

0.427

0.698

0.639

0.836

0.585

20

0.217

0.641

0.858

0.814

1.068

0.904

25

0.284

0.820

1.075

1.174

1.300

30

0.423

1.006

1.378

1.526

1.188 -

35

0.512

1.160

1.787

1.783

-

40

0.605

1.457

1.972

2.159

-

1.806 -

45

0.777

1.882

2.325

-

-

50

0.993

2.113

-

-

55

1.239

2.203

2.764 -

2.852 -

-

-

60

1.526

2.770

-

-

-

-

129

1.468

Appendix B Table B-11: The relationship between the apparent rate constant and the temperature of LGSF dye solution using commercial nano TiO2. T (K)

k (min–1)

283.15 288.15 293.15 298.15 303.15 308.15

0.0274 0.0400 0.0505 0.0539 0.0898 0.0491

Table B-12: The variation of 1/T with ln(k). (103/T) (K)

ln(k) (min–1)

3.53 3.47 3.41 3.35 3.29 3.24

-3.597 -3.219 -2.986 -2.921 -2.410 -3.014

Table B-13: The variation of 1/T with ln(k/T). (103/T) (K)


3.53 3.47 3.41 3.35 3.29 3.24

-9.243 -8.882 -8.666 -8.618 -8.124 -8.744

130

Appendix B Table B-14: Thermodynamic parameters of the decolorization of LGSF dye using commercial nano TiO2. Ea (kJ mol–1)


37.032

34.667

∆S‡ (kJ mol–1 K–1) ∆G‡288.15 (kJ mol–1) – 0.152

80.664

Table B-15: The relationship between PDE and irradiation time for different temperatures using commercial nano TiO2. t (min)

PDE for different temperatures (%) 283.15

288.15

293.15

298.15

303.15

308.15

0

0.000

0.000

0.000

0.000

0.000

0.000

5

5.413

10.149

18.696

18.676

15.625

18.838

10

11.491

24.627

36.957

36.842

36.719

33.267

15

15.480

34.776

50.217

47.199

56.641

44.289

20

19.468

47.313

57.609

55.688

65.625

59.519

25

24.691

55.970

65.870

69.100

69.531

72.745

30

34.473

63.433

74.783

78.268

83.203

76.954

35

40.076

68.657

83.261

83.192

-

83.567

40

45.394

76.716

86.087

88.455

-

-

45

54.036

84.776

90.217

94.228

-

-

50

62.963

87.910

93.696

-

-

-

55

71.035

88.955

-

-

-

-

60

78.253

93.731

-

-

-

-

131

Appendix C Table C-1: The change of ln(C0/Ct) with irradiation time for different mixing ratios of the prepared TiO2 calcinated at 600 and 800 °C.

t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

1:4 0.000 0.105 0.280 0.448 0.973 1.282 1.517 1.825 2.234 -

ln(C0/Ct) for different mixing ratios 600 °C 800 °C 3:10 1:10 1:4 3:10 0.000 0.000 0.000 0.000 0.612 0.346 0.076 0.075 1.115 0.673 0.201 0.136 1.687 1.052 0.287 0.219 2.366 1.399 0.466 0.285 2.943 1.827 0.595 0.351 2.392 0.787 0.458 0.898 0.515 1.108 0.607 1.398 0.704 1.646 0.775 2.096 0.864 2.391 0.949

1:10 0.000 0.081 0.142 0.206 0.296 0.382 0.481 0.526 0.645 0.836 0.890 0.970 1.112

Table C-2: The relationship between the apparent rate constant and the mixing ratio for the prepared TiO2 calcinated at 600 and 800 °C. k (min–1)

v/v 1:4 3:10 1:10

600 °C 0.0518 0.1167 0.0746

132

800 °C 0.0329 0.0162 0.0205

Appendix C Table C-3: The relationship between PDE and irradiation time for different mixing ratios of the prepared TiO2 calcinated at 600 and 800 °C.

t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

1:4 0.000 9.935 24.387 36.129 62.194 72.258 78.065 83.871 89.290 -

PDE for different mixing ratios (%) 600 °C 800 °C 3:10 1:10 1:4 3:10 0.000 0.000 0.000 0.000 45.758 29.250 7.349 7.231 67.224 48.995 18.193 12.750 81.491 65.082 24.940 19.696 90.617 75.320 37.229 24.833 94.730 83.912 44.819 29.591 90.859 54.458 36.727 59.277 40.247 66.988 45.480 75.301 50.523 80.723 53.949 87.711 57.850 90.843 61.275

1:10 0.000 7.761 13.258 18.593 25.627 31.770 38.157 40.905 47.534 56.669 58.933 62.086 67.098

Table C-4: The change of ln(C0/Ct) with irradiation time for different concentrations of LGSF dye solution using prepared TiO2. t (min) 0 5 10 15 20 25 30 35

ln(C0/Ct) for different concentrations of dye solution (ppm) 15 20 25 0.000 0.000 0.000 0.323 0.612 0.302 0.662 1.115 0.724 0.945 1.687 1.028 1.374 2.366 1.505 1.683 2.943 1.773 2.173 2.801

133

Appendix C Table C-5: The relationship between the apparent rate constant and the initial concentration of LGSF dye solutions using prepared TiO2. LGSF conc. (ppm) 15 20 25

k (min–1) 0.0737 0.1167 0.0748

Table C-6: The relationship between PDE and irradiation time for different initial concentrations of LGSF dye solution using prepared TiO2. t (min) 0 5 10 15 20 25 30 35

PDE for different concentrations of dye solution (%) 15ppm 20ppm 25ppm 0.000 0.000 0.000 27.578 38.574 26.091 48.441 67.540 51.518 61.151 77.584 64.231 74.700 87.045 77.799 81.415 91.703 83.017 88.615 93.928

134

Appendix C

Table C-7: The change of ln(C0/Ct) with irradiation time for different masses of prepared TiO2. t (min)

0 5 10 15 20 25 30 35 40 45 50 55 60

0.05 0.000 0.081 0.159 0.253 0.360 0.476 0.593 0.712 0.849 0.983 1.134 1.317 1.482

ln(C0/Ct) for different masses (g) 0.1 0.15 0.2 0.25 0.3 0.000 0.000 0.000 0.000 0.000 0.196 0.404 0.612 0.275 0.220 0.383 0.821 1.115 0.745 0.630 0.590 1.128 1.687 1.120 0.812 0.762 1.402 2.366 1.761 1.016 0.951 1.715 2.943 2.287 1.360 1.213 2.136 1.597 1.495 2.703 1.980 1.826 2.154 2.461 2.745 -

0.4 0.000 0.126 0.363 0.570 0.884 1.120 1.307 1.580 1.875 2.111 2.562 -

Table C-8: The relationship between the apparent rate constant and the mass of the prepared TiO2. Catalyst dosage (g)

k (min–1)

0.05 0.1 0.15 0.2 0.25 0.3 0.4

0.0231 0.0463 0.0734 0.1167 0.0858 0.0546 0.0469

135

Appendix C Table C-9: The relationship between PDE and irradiation time for different masses of the prepared TiO2. t (min)

0 5 10 15 20 25 30 35 40 45 50 55 60

0.05 0.000 7.781 14.677 22.370 30.239 37.843 44.739 50.928 57.206 62.599 67.816 73.210 77.277

PDE for different masses (%) 0.1 0.15 0.2 0.25 0.3 0.000 0.000 0.000 0.000 0.000 17.833 33.245 45.758 24.023 19.734 31.831 55.996 67.224 52.539 46.728 44.583 67.637 81.491 67.383 55.624 53.308 75.397 90.617 82.813 63.804 61.361 82.011 94.730 89.844 74.335 70.278 88.183 79.755 77.565 93.298 86.196 83.893 88.399 91.467 93.576 -

-

-

-

-

0.4 0.000 11.816 30.469 43.457 58.691 67.383 72.949 79.395 84.668 87.891 -

Table C-10: The change of ln(C0/Ct) with irradiation time for different initial pH value of the dye solution in the presence of prepared TiO2 calcinated at 600 °C. t (min) 0 5 10 15 20 25 30 35 40 45 50

ln(C0/Ct) at different pH 3 0.000 0.341 0.676 1.112 1.597 2.097 2.233 2.925 -

4 0.000 0.435 0.874 1.270 1.672 2.083 2.619 -

5 0.000 0.426 1.002 1.541 2.127 2.400 -

6 0.000 0.464 1.127 1.493 2.182 2.913 -

136

7.3 0.000 0.612 1.115 1.687 2.366 2.943 -

8 0.000 0.426 0.682 1.057 1.353 1.622 2.189 2.460 2.654 -

9 0.000 0.134 0.371 0.551 0.736 0.945 1.246 1.455 1.634 2.085 2.242

Appendix C Table C-11: The relationship between the apparent rate constant and the initial pH of LGSF dye solution in the presence of prepared TiO2 calcinated at 600 °C. pH

k (min–1)

3 4 5 6 7.3 8 9

0.0797 0.0853 0.1002 0.1109 0.1167 0.0697 0.0427

Table C-12: The relationship between PDE and irradiation time for different initial pH of LGSF dye solution prepared TiO2 calcinated at 600 °C. PDE for different pH (%)

t (min)

3

4

5

6

7.3

8

9

0

0.000

0.000

0.000

0.000

0.000

0.000

0.000

5

28.896 35.269 34.705 37.095 45.758 34.673 12.572

10

49.147 58.289 63.291 67.607 67.224 49.462 30.997

15

67.122 71.911 78.586 77.534 81.491 65.255 42.341

20

79.750 81.204 88.080 88.715 90.617 74.156 52.095

25

87.713 87.540 90.928 94.566 94.730 80.258 61.127

30

89.280 92.714

-

-

-

88.801 71.243

35

94.631

-

-

-

-

91.457 76.662

40

-

-

-

-

-

92.965 80.491

137

Appendix C Table C-13: The change of ln(C0/Ct) with irradiation time for different temperatures in the presence of prepared TiO2 calcinated at 600 °C. t (min) 0 5 10 15 20 25 30 35 40 45 50 55 60

ln(C0/Ct) for different temperatures (K) 283.15

288.15

293.15

0.000 0.060 0.181 0.302 0.362 0.476 0.617 0.729 0.866 1.065 1.177 1.463 1.698

0.000 0.183 0.499 0.779 1.026 1.277 1.364 1.715 2.117 2.470 2.713 3.188 3.772

0.000 0.336 0.626 0.965 1.359 1.792 2.128 2.475 3.138 -

298.15 0.000 0.612 1.115 1.687 2.366 2.943 -

303.15 0.000 0.612 1.115 1.687 2.366 2.943 -

308.15 0.000 0.299 0.633 0.928 1.377 1.807 2.137 2.655 3.077 -

Table C-14: The relationship between the apparent rate constant and the temperature of LGSF dye solution using prepared TiO2 calcinated at 600 °C. T (K)

k (min–1)

283.15 288.15 293.15 298.15 303.15 308.15

0.0267 0.0556 0.0726 0.0767 0.1167 0.0734

138

Appendix C Table C-15: The variation of 1/T with ln(k). (103/T) (K)

ln(k) (min–1)

3.53 3.47 3.41 3.35 3.29 3.24

-3.623 -2.890 -2.623 -2.568 -2.148 -2.612

Table C-16: The variation of 1/T with ln(k/T). (103/T) (K)


3.53 3.47 3.41 3.35 3.29 3.24

-9.269 -8.553 -8.303 -8.265 -7.862 -8.342

Table C-17: Thermodynamic parameters of the decolorization of LGSF dye using prepared TiO2 nanoparticles calcinated at 600 °C. Ea (kJ mol–1)


45.780

42.969

∆S‡ (kJ mol–1 K–1) ∆G‡288.15 (kJ mol–1) -0.120

139

79.707

Appendix C Table C-18: The relationship between PDE and irradiation time for different temperatures using prepared TiO2 nanoparticles calcinated at 600 °C. t (min)

PDE for different temperatures (%) 283.15

288.15

293.15

298.15

303.15

308.15

0

0.000

0.000

0.000

0.000

0.000

0.000

5

5.855

16.712

28.571

45.758

28.348

25.869

10

16.567

39.310

46.514

67.224

48.261

46.888

15

26.081

54.127

61.905

81.491

63.130

60.469

20

30.339

64.141

74.320

90.617

79.043

74.778

25

37.858

72.124

83.333

94.730

83.826

83.589

30

46.041

74.425

88.095

-

91.652

88.197

35

51.763

82.003

91.582

-

92.967

40

57.951

87.957

95.663

-

96.348 -

45

65.536

91.543

-

-

-

-

50

69.195

93.369

-

-

-

-

55

76.846

95.873

-

-

-

-

60

81.703

97.700

-

-

-

-

140

95.392

‫\‪Ïê˜£‬‬ ‫ﺗﺘﻀﻤﻦ ھﺬه اﻟﺪراﺳﺔ ﺛﻼﺛﺔ أﺟﺰاء‪:‬‬ ‫اﻟﺠﺰء اﻷول‪ :‬ﺗﻀﻤﻦ ﺗﺤﻀﯿﺮ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم ﺑﻄﺮﯾﻘﺔ اﻟﺼﻮل ‪ -‬ﺟﻞ وﺑﻨﺴﺐ ﻣﺰج ﻣﺨﺘﻠﻔﺔ )‪،3:10 ،1:10‬‬ ‫‪ (1:4‬ﺑﯿﻦ رﺑﺎﻋﻲ ﻛﻠﻮرﯾﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم واﻟﻜﺤﻮل اﻷﺛﯿﻠﻲ‪ ،‬وﺗﻠﺪﯾﻦ اﻟﻨﻮاﺗﺞ ﺑﺪرﺟﺘﻲ ‪ 600‬و ‪ 800‬م‪.°‬‬ ‫اﻟﺠﺰء اﻟﺜﺎﻧﻲ‪ :‬وﺗﻀﻤﻦ ھﺬا اﻟﺠﺰء دراﺳﺔ ﺧﻮاص أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري واﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي واﻟﻤﺤﻀﺮ‬ ‫ﻓﻲ ھﺬه اﻟﺪراﺳﺔ ﻣﻦ ﺧﻼل ﺗﻘﻨﯿﺔ ﺣﯿﻮد اﻷﺷﻌﺔ اﻟﺴﯿﻨﯿﺔ وﻣﺠﮭﺮ اﻟﻘﻮة اﻟﺬرﯾﺔ واﻟﻔﻠﻮرة‪.‬‬ ‫ﺑﯿﻨﺖ أطﯿﺎف اﻷﺷﻌﺔ اﻟﺴﯿﻨﯿﺔ ﻋﺪم وﺟﻮد طﻮر اﻟﺮوﺗﺎﯾﻞ واﻟﺒﺮوﻛﺎﯾﺖ ﻓﻲ اﻟﻨﻤﺎذج اﻟﻤﻠﺪﻧﺔ ﺗﺤﺖ درﺟﺔ ﺣﺮارة‬ ‫‪ °600‬ﻣﺌﻮﯾﺔ‪ ،‬ﻓﻲ ﺣﯿﻦ أن طﻮر اﻟﺮوﺗﺎﯾﻞ ﺑﺪأ ﺑﺎﻟﻈﮭﻮر ﻓﻲ اﻟﻨﻤﺎذج اﻟﻤﻠﺪﻧﺔ ﺗﺤﺖ درﺟﺔ ﺣﺮارة ‪ °800‬ﻣﺌﻮﯾﺔ‪.‬‬ ‫ﺗﻢ ﺣﺴﺎب ﻣﻌﺪل اﻟﺤﺠﻮم اﻟﺒﻠﻮرﯾﺔ واﻟﺤﺠﻮم اﻟﺒﻠﻮرﯾﺔ ﻷﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري واﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي‬ ‫واﻟﻤﺤﻀﺮ ﺑﻮاﺳﻄﺔ ﻣﻌﺎدﻟﺔ ﺷﯿﺮر وﻣﻌﺎدﻟﺔ ﺷﯿﺮر اﻟﻤﻌﺪﻟﺔ‪.‬‬ ‫ﺑﯿﻨﺖ ﺻﻮر ﻣﺠﮭﺮ اﻟﻘﻮة اﻟﺬرﯾﺔ أن ﻣﻌﻈﻢ أﺷﻜﺎل أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮة ھﻲ ﻛﺮوﯾﺔ‪.‬‬ ‫ﻛﻤﺎ ﺑﯿﻨﺖ أطﯿﺎف اﻟﻔﻠﻮرة أن طﺎﻗﺔ ﻓﺠﻮة اﻟﺤﺰﻣﺔ ﻟﻠﻨﻤﺎذج اﻟﻤﺤﻀﺮة ﺗﺮواﺣﺖ ﺑﯿﻦ )‪ 3.56 – 3.51‬اﻟﻜﺘﺮون‬ ‫ﻓﻮﻟﺖ( وھﺬه اﻟﻘﯿﻤﺔ أﻛﺒﺮ ﺑﻘﻠﯿﻞ ﻣﻦ اﻟﻘﯿﻢ اﻟﻤﺴﺠﻠﺔ‪.‬‬ ‫اﻟﺠﺰء اﻟﺜﺎﻟﺚ‪ :‬ﺗﻀﻤﻦ ھﺬا اﻟﺠﺰء دراﺳﺔ ﺗﺄﺛﯿﺮ درﺟﺔ ﺣﺮارة اﻟﺘﻠﺪﯾﻦ ﻋﻠﻰ اﻟﻔﻌﺎﻟﯿﺔ اﻟﻀﻮﺋﯿﺔ ﻟﻨﻤﺎذج أﻛﺴﯿﺪ‬ ‫اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮة‪ ،‬وﻗﺪ ﺑﯿﻨﺖ اﻟﻨﺘﺎﺋﺞ أن أﻓﻀﻞ ﻧﺴﺒﺔ ﻣﺰج ھﻲ )‪ (3:10‬اﻟﻤﻠﺪن ﺗﺤﺖ درﺟﺔ ﺣﺮارة ‪°600‬‬ ‫ﻣﺌﻮﯾﺔ‪.‬‬ ‫ﺗﻀﻤﻦ ھﺬا اﻟﺠﺰء ﻛﺬﻟﻚ دراﺳﺔ اﻟﻌﻮاﻣﻞ اﻟﻤﺨﺘﻠﻔﺔ اﻟﺘﻲ ﺗﺆﺛﺮ ﻋﻠﻰ ﻋﻤﻠﯿﺔ اﻟﺘﺤﻄﻢ اﻟﻀﻮﺋﻲ اﻟﻤﺤﻔﺰ ﻟﺼﺒﻐﺔ‬ ‫‪ Light Green SF Yellowish‬ﻣﻊ وﺟﻮد أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري واﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي واﻟﻨﺎﻧﻮي اﻟﻤﺤﻀﺮ‪،‬‬ ‫وﺗﺸﻤﻞ ھﺬه اﻟﻌﻮاﻣﻞ ﺗﺮﻛﯿﺰ اﻟﺼﺒﻐﺔ وﻛﻤﯿﺔ اﻟﻌﺎﻣﻞ اﻟﻤﺴﺎﻋﺪ‪ ،‬واﻟﺪاﻟﺔ اﻟﺤﺎﻣﻀﯿﺔ اﻻﺑﺘﺪاﺋﯿﺔ ﻟﻠﻤﺤﻠﻮل ودرﺟﺔ‬ ‫اﻟﺤﺮارة‪.‬‬ ‫ﺗﻤﺖ دراﺳﺔ ﺗﺄﺛﯿﺮ اﻟﺘﺮﻛﯿﺰ اﻻﺑﺘﺪاﺋﻲ ﻟﻠﺼﺒﻐﺔ ﺑﺎﺳﺘﺨﺪام ﺗﺮاﻛﯿﺰ ﻣﺨﺘﻠﻔﺔ )‪ (ppm 50-20‬وﻗﺪ وﺟﺪ أن اﻟﺘﻔﺎﻋﻞ‬ ‫ﯾﺨﻀﻊ ﻟﻠﻤﺮﺗﺒﺔ اﻷوﻟﻰ اﻟﻜﺎذﺑﺔ‪.‬‬

‫أ‬

‫ﻋﯿﻨﺖ ﻛﻤﯿﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺜﻠﻰ وﻟﻜﻞ ﻧﻮع ﻣﻦ أﻧﻮاع اﻟﺘﯿﺘﺎﻧﯿﻮم‪ ،‬وﻛﺎﻧﺖ أﻓﻀﻞ ﻗﯿﻤﺔ ﺑﺎﻟﻨﺴﺒﺔ ﻷﻛﺴﯿﺪ‬ ‫اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري )‪ 0.3‬ﻏﻢ( ﻓﻲ ‪ 100‬ﻣﻠﻠﺘﺮ‪ ،‬أﻣﺎ ﺑﺎﻟﻨﺴﺒﺔ ﻷﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻨﺎﻧﻮي اﻟﺘﺠﺎري واﻟﻤﺤﻀﺮ ﻓﻘﺪ‬ ‫ﻛﺎﻧﺖ )‪ 0.2‬ﻏﻢ( ﻓﻲ ‪ 100‬ﻣﻠﻠﺘﺮ‪.‬‬ ‫أﻣﺎ ﺑﺎﻟﻨﺴﺒﺔ إﻟﻰ اﻟﺪاﻟﺔ اﻟﺤﺎﻣﻀﯿﺔ اﻻﺑﺘﺪاﺋﯿﺔ ﻟﻤﺤﺎﻟﯿﻞ اﻟﺼﺒﻐﺔ ﻓﻘﺪ ﺑﯿﻨﺖ اﻟﻨﺘﺎﺋﺞ أن أﻓﻀﻞ داﻟﺔ ﺣﺎﻣﻀﯿﺔ اﺑﺘﺪاﺋﯿﺔ‬ ‫ﺗﺴﺎوي )‪ (7.3‬ﻟﺠﻤﯿﻊ أﻛﺎﺳﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺴﺘﺨﺪﻣﺔ ﻓﻲ ھﺬا اﻟﻌﻤﻞ‪.‬‬ ‫ﺗﻤﺖ ﻛﺬﻟﻚ دراﺳﺔ ﺗﺄﺛﯿﺮ درﺟﺔ اﻟﺤﺮارة ﺑﺎﻻﺳﺘﻌﺎﻧﺔ ﺑﻤﻌﺎدﻟﺔ أرﯾﻨﻮس‪ ،‬وﻗﺪ ﺑﯿﻨﺖ اﻟﻨﺘﺎﺋﺞ أن ﺳﺮﻋﺔ اﻟﺘﻔﺎﻋﻞ‬ ‫ﺗﺰداد ﺑﺎزدﯾﺎد درﺟﺔ اﻟﺤﺮارة ﻣﻦ )‪ (278.15-293.15‬ﻛﻠﻔﻦ‪ ،‬وھﺬا ﯾﺪل ﻋﻠﻰ أن ﺗﻔﺎﻋﻞ اﻹزاﻟﺔ اﻟﻠﻮﻧﯿﺔ‬ ‫اﻟﻀﻮﺋﯿﺔ ﻟﺼﺒﻐﺔ ‪ Light Green‬ھﻮ ﺗﻔﺎﻋﻞ ﻣﺎص ﻟﻠﺤﺮارة‪ .‬ﺗﻢ ﺣﺴﺎب طﺎﻗﺔ اﻟﺘﻨﺸﯿﻂ ﻟﻜﻞ ﻧﻮع ﻣﻦ أﻧﻮاع‬ ‫أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺴﺘﺨﺪﻣﺔ ﻓﻲ اﻟﻌﻤﻞ‪ ،‬وﻗﺪ وﺟﺪ أﻧّﮭﺎ ﺗﺴﺎوي )‪ (48.676‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ‬ ‫اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري‪ ،‬و )‪ (37.032‬ﻛﯿﻠﻮﺟﻮل\ ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي‪ ،‬و)‪(45.780‬‬ ‫ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮ ﺑﻨﺴﺒﺔ ‪ 1:3‬واﻟﻤﺤﺮوق ﺑﺪرﺟﺔ ‪ 600‬م‪.°‬‬ ‫ﻛﻤﺎ ﺗﻢ ﺣﺴﺎب ﻗﯿﻢ اﻹﻧﺘﺮوﺑﯿﺔ واﻹﻧﺜﺎﻟﺒﯿﺔ ﺑﺎﺳﺘﺨﺪام ﻣﻌﺎدﻟﺔ إﯾﺮﻧﺞ – ﺑﻮﻻﻧﻲ‪ ،‬وﻗﺪ وﺟﺪ أن ﻗﯿﻢ اﻹﻧﺘﺮوﺑﯿﺔ‬ ‫ﺗﺴﺎوي )‪ (– 0.107‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري‪ ،‬و)‪ (– 0.152‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ‬ ‫أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي‪ ،‬و)‪ (– 0.120‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮ‪ ،‬وﻗﺪ وﺟﺪ‬ ‫ﺑﺄن ﻗﯿﻢ اﻹﻧﺘﺮوﺑﯿﺔ ﻣﻨﺨﻔﻀﺔ ﻧﺘﯿﺠﺔ ﻟﻨﻘﺼﺎن اﻟﻌﺸﻮاﺋﯿﺔ‪.‬‬ ‫أﻣﺎ ﻗﯿﻢ اﻹﻧﺜﺎﻟﺒﯿﺔ ﻓﻘﺪ وﺟﺪت ﺑﺄﻧﮭﺎ ﺗﺴﺎوي )‪ (46.292‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري‪،‬‬ ‫و)‪ (34.667‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي‪ ،‬و)‪ (43.416‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل‪ ،‬ﻓﻲ‬ ‫ﻣﺎص ﻟﻠﺤﺮارة‪.‬‬ ‫ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮ‪ ،‬وﺑﯿّﻨﺖ اﻟﻨﺘﺎﺋﺞ أن اﻟﺘﻔﺎﻋﻞ‬ ‫ّ‬ ‫ﺗ ّﻢ ﺣﺴﺎب اﻟﻄﺎﻗﺔ اﻟﺤﺮة ﻟﻜﻞ ﻧﻮع ﻣﻦ أﻧﻮاع أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺴﺘﺨﺪﻣﺔ ﻓﻲ اﻟﻌﻤﻞ‪ ،‬وﻗﺪ وﺟﺪ أﻧّﮭﺎ ﺗﺴﺎوي‬ ‫)‪ (78.847‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري‪ ،‬و )‪ (80.664‬ﻛﯿﻠﻮﺟﻮل\ ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ‬ ‫اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري اﻟﻨﺎﻧﻮي‪ ،‬و)‪ (79.707‬ﻛﯿﻠﻮﺟﻮل\ﻣﻮل ﻓﻲ ﺣﺎﻟﺔ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﻤﺤﻀﺮ‪ ،‬وﺑﯿّﻨﺖ اﻟﻨﺘﺎﺋﺞ‬ ‫ﺑﺄن اﻟﺘﻔﺎﻋﻞ ﻏﯿﺮ ﺗﻠﻘﺎﺋﻲ‪.‬‬

‫ب‬

‫ﺟﻤﮭﻮرﯾﺔ اﻟﻌﺮاق‬ ‫وزارة اﻟﺘﻌﻠﯿﻢ اﻟﻌﺎﻟﻲ واﻟﺒﺤﺚ اﻟﻌﻠﻤﻲ‬ ‫ﺟﺎﻣﻌﺔ اﻟﻜﻮﻓﺔ ‪ -‬ﻛﻠﯿﺔ اﻟﻌﻠﻮم‬ ‫ﻗﺴﻢ اﻟﻜﯿﻤﯿﺎء‬

‫اﻟﺘﻔﻜﻚ اﻟﻀﻮﺋﻲ ﻟﺼﺒﻐﺔ ‪Light Green‬‬ ‫ﺑﺎﺳﺘﺨﺪام ﺛﻨﺎﺋﻲ أﻛﺴﯿﺪ اﻟﺘﯿﺘﺎﻧﯿﻮم اﻟﺘﺠﺎري واﻟﻨﺎﻧﻮي‬ ‫ﻛﻌﺎﻣﻞ ﻣﺴﺎﻋﺪ‬ ‫رﺳﺎﻟﺔ ﻣﻘﺪﻣﺔ إﻟﻰ‬ ‫ﻣﺠﻠﺲ ﻛﻠﯿﺔ اﻟﻌﻠﻮم ‪ /‬ﺟﺎﻣﻌﺔ اﻟﻜﻮﻓﺔ وھﻲ ﺟﺰء ﻣﻦ ﻣﺘﻄﻠﺒﺎت ﻧﯿﻞ ﺷﮭﺎدة‬ ‫اﻟﻤﺎﺟﺴﺘﯿﺮ ﻓﻲ اﻟﻜﯿﻤﯿﺎء‬ ‫ﺗﻘﺪم ﺑﮭﺎ‬ ‫ﻣﺤﻤﺪ طﺎھﺮ ﻋﯿﺴﻰ‬ ‫ﺑﻜﻠﻮرﯾﻮس ‪) 2012‬ﺟﺎﻣﻌﺔ ﻛﺮﺑﻼء(‬

‫ﺑﺈﺷﺮاف‬ ‫أ‪.‬م‪.‬د‪ .‬ﻋﺎﻣﺮ ﻣﻮﺳﻰ ﺟﻮدة‬ ‫‪ 1438‬ھـ‬

‫أ‪.‬م‪.‬د‪ .‬ﻟﻤﻰ ﻣﺠﯿﺪ أﺣﻤﺪ‬ ‫‪ 2016‬م‬