ii

BIOSORPTIVE POTENTIAL AND BINDING

SITES OF HEAVY METALS ON THE CELL

WALL OF SELECTED FILAMENTOUS

ALGAE

Submitted to The GC University, Lahore, Pakistan

In partial fulfillment of the requirements For the award of the Degree of

DOCTOR OF PHILOSOPHY

IN

ZOOLOGY By

ATIF YAQUB

Session: 2005-2008

Roll

No:

2 4 3 P H D Z 0 5

DEPARTMENT OF ZOOLOGY

GC UNIVERSITY, LAHORE

iii

Dedicated

To

MY FAMILY

MY WIFE, NAUREEN

AND MY SONS,

HASSAAN AND

REHAN

Whose continuous support and patience have

enabled me to accomplish my overly long academic pursuit

iv

RESEARCH COMPLETION CERTIFICATE

Certified that the research work contained in this thesis entitled

“Biosorptove Potential and Binding Sites of Heavy Metals on the Cell Wall of

Selected Filamentous Green Algae” has been carried out and completed by

Mr. Atif Yaqub, Roll No. 243-GCU-PHD-Z-05 under my supervision and is

being submitted for the award of PhD degree in Zoology research work.

Research Supervisor

Prof. Dr. Sharif Mughal

CHAIRPERSON, DEPARTMENT OF ZOOLOGY

GOVERNMENT COLLEGE

UNIVERSITY,

LAHORE

Dated:

Submitted through

Prof. Dr. Sharif Mughal

CHAIPERSON, DEPARTMENT OF ZOOLOGY

GOVERNMENT COLLEGE UNIVERSITY,

LAHORE

Controller of Examination

Government College University, Lahore

v

IN THE NAME OF

ALLAH

The Most Merciful

The Most Beneficent

The Most Knowing

“It is ALLAH who has made for

you the earth as a resting place and

the sky as a canopy and has given you

shape and made your shapes beautiful

and has provided for you sustenance of

things pure and good. Such is ALLAH

your Lord. So glory to ALLAH.

The Lord of the world”

(Surah 40 Ghafir:64)

vi

GOLDEN SAYINGS OF THE HOLY PROPHET

HAZRAT MUHAMMAD (Peace be upon him)

“Keep your thought well composed, and search

the fact of wisdom for it, otherwise mind

gets weary and the people get weary,

so ponder into the knowledge and

science thought fully and

search for new facts

and ideas”

vii

D E C L A R A T I O N

I, Atif Yaqub, Roll No. 243-GCU-PHD-Z05 student of PhD in the

subject of Zoology Session 2005-2008 hereby declare that the matter

printed in this thesis, entitled “Biosorptive Potential and Binding Sites of

Heavy Metals on the Cell Wall of Selected Filamentous Green Algae” is my

own work and has not been printed, published and submitted as a

research work, thesis or publication in any University, Research

Institution, etc. in Pakistan or abroad.

Dated: Signature of Deponent

viii

ACKNOWLEDGMENTS

All praises to Almighty, Allah, the Lord. Then, I would like to gratefully acknowledge

the many people whose support made it possible for me to complete this research. I would like to

pay my deep gratitude to my supervisor; Prof. Dr. Muhammad Sharif Mughal, Professor and

Chairperson at Department of Zoology, GC University, Lahore, whose sincere efforts to provide

me counseling to his level best have made this project possible. I am permanently indebted to

professor at Aquatic Entomology and Ecology laboratory, Mid-Florida Research and Education

Centre, University of Florida whose moral and practical support has always been with me; and

who has taught me much about good lab practice, experimental design, and, most of all, the

required self-discipline. Dr. Muhammad Ayub is worth to pay gratitude for sharing his

knowledge, encouragement and advice which has helped in my graduate studies. I would like to

pay special thanks to Dr. Ahmad Adnan, Associate Professor at Department of Chemistry, GC

University, Lahore, for taking his time to teach me various procedures and sharing his extensive

knowledge of chemistry. I would also like to extend special thanks to the members of research

committee of the Department of Zoology, GC University, Lahore, Dr. Azizullah, Dr. M.Sharif

Mughal, Dr. Nusrat Jahan and Dr. Zaheer Ahmad, whose valuable suggestions helped me to

develop a better research thesis.

I would acknowledge the cooperation of my friends, Syed Umer Farooq Rizvi and Waqar

Nasir, PhD scholars at Chemistry department at University of the Punjab, Lahore for their

technical help and guidance. I also gratefully acknowledge Dr. Sadanand Dhekeney, Visiting

Assistant Professor at University of Florida, USA and Dr. Abhijit Mazumdar, Associate

Professor of Zoology at University of Burdwan, West Bengal, India, for their pleasant company

and continuous guidance during my research work at University of Florida. Spending time and

sharing some special moments with them has been a wonderful experience for me.

I would also like to extend special thanks to my friends, Chand Raza and Dr. Khalid

Anjum, faculty members of department of Zoology, GCU, Lahore, who have always stood by

me and left no stone unturned to help me as I progressed through my studies; finding their time

ix

to share cups of tea has been appreciable in particular. A very special and outreaching help from

Muahmmad Faizan Naeem, a B.Sc.(H) research student of the same department is also

admirable. Other faculty members of this department have also been considerate and helpful

specially Dr. Nazish Mazhar Ali and Imran Sohail. All the laboratory and office staff of the

Department of Zoology also deserves thanks for their cooperation during this project.

I am permanently indebted to my family who have patiently supported my overly long

academic pursuits. I would specifically like to thank my wife, Naureen Mumtaz for extending

her commitments and dedications which have helped me to accomplish this project; my brother

Asim Yaqub, who has been ready to help whenever I fell in need; my father, Muhammad Yaqub,

who has laid my foundations and has ever been a source of inspiration for me; my mother, Talat

Yaqub; and my parents in law, Mumtaz Anwar and Kalsoom whose prayers have always been

with me.

My sons, Hassaan Atif and Rehan Atif also deserve gratitude because they are the reason

of my being in action.

Also thanks to many more people who have helped me one way or the other during my

studies.

Last, but the not the least, I would like to pay special thanks to Higher Education

Commission (HEC), Government of Pakistan, for their support in this project. HEC has provided

me opportunity to conduct my research at University of Florida, for six months, resulting in

generation of some important research data and a research publication, under the program,

International Research Support Initiative Program in 2007.

x

ABSTRACT

Indiscriminate release of heavy metal pollutants into the environment from point

and non-point industrial sources has posed a major threat to all kinds of organisms

inhabiting aquatic and terrestrial habitats. The application of biosorption i.e., removal of

heavy metal ions by the use of biomass, has emerged as a promising technique in the

past few years. Utilization of green filamentous algae in this technology still remains

largely unexplored. In the present study, the biosorption capacities of biomass of

filamentous green algae, Spirogyra cummunis, Cladophora delmatica, and Spirogyra

spp. were evaluated for toxic heavy metals, such as Cadmium, Cd (II) and hexavalent

Chromium, Cr (VI). The biosorptive binding sites were studied with the help of Scanning

Electron Microscopy (SEM) and Fourier Transform Infra-red Spectrometer (FTIR).

Results revealed that the rate and extent of uptake were influenced by pH, contact time,

and biosorbent concentration. The optimum pH value for uptake of Cd (II) was found to

be 5.0 and that of Cr (VI) 4.0 by all the studied biosorbents. The equilibrium sorption

data for Cd (II) at pH 5.0 and that of Cd (II) at pH 4.0 were described by various

adsorption isotherms, such as Langmuir, Freundlich, and Temkin models. Langmuir

isotherm was found to be the best suited for the interpretation of acquired data, showing

monolayer adsorption and Freundlich theorem, the worst. Values of Cd (II) sorption

capacity, (qmax) for the studied species were found to be 1.44, 11.9 and 14.42 and those

of Cr (VI) were 498, 411 and 312, respectively. Kinetic and thermodynamic parameters

were also studied. The results showed that pseudo-second order kinetics was suitable

for the interpretation of data and thermodynamically biosorption was found to be

feasible and spontaneous under the given conditions, in case of all the biosorption

investigations undertaken in the present study. SEM and FTIR revealed biosorptive

binding sites and possible electronegative functional groups, such as carboxyl, hydroxyl,

carbonyl, etc., on the surface of biosorbents which could favor the binding of cations,

such as Cd (II) and Cr (VI) ions.

xi

The results of the present study indicate filamentous green algae as a potential

candidate for their application in biosorption at industrial scale. Its optimization for

industry scale usage and hybridization with already existing processes may possibly be

realized in improving wastewater treatment reactor, and therefore the corresponding

novel reactors may be designed which will be cost-effective as well.

Keywords: Biosorption, Algae, Heavy metals, Wastewater, Equilibrium, Kinetics,

Thermodynamics, FTIR, SEM.

xii

TABLE OF CONTENTS

CHAPTER 1

INTRODUCTION ······················································································································· 1

CHAPTER 2

LITERATURE REVIEW ······································································································11

CHAPTER 3

MATERIALS AND METHODS ················································································································ 23

3.1. MATERIALS ········································································ 23

3.1.1. BIOSORBENTS ······························································· 23

3.1.1.1. Biosorbent I: ········································································ 23

3.1.1.2. Biosorbent II: ······································································· 23

3.1.1.3. Biosorbent III: ······································································ 23

3.1.1.4. Culture conditions of Spirogyra communis ·································· 23

3.1.1.5. Pretreatment of Biomass ························································ 24

3.1.2- PREPARATION OF HEAVY METAL (Cd II AND

Cr VI) SOLUTIONS ···································································· 24

3.1.2.1 Cadmium Solution ·································································· 24

3.1.2.2 Chromium Solution ································································· 24

3.2 METHODS AND PROCEDURES ·················································· 25

3.2.1 BISORPTION STUDIES ······························································ 25

3.2.1.1 Glassware and Apparatus ······················································ 25

3.2.1.2. Biosorption evaulation ……………………………………………………25

3.2.1.3. Transport of Samples for Analysis ············································ 26

3.2.1.4. Analysis by Atomic Absorption Spectrometer ······························ 26

3.2.2 EVALUATION OF BIOSORPTIVE BINDING

SITES ············································································ 26

3.2.2.1 Fourier Transform Infra red Spectroscopy (FTIR) ························ 26

3.2.2.2 Scanning Electron Microscopy (SEM) ······································· 27

3.2.3. MATHEMATICAL MODELING AND

INTERPRETATION OF DATA ··················································· 27

xiii

3.2.3.1. Equilibrium Isotherm Modeling ·················································· 27

3.2.3.2. Kinetic Studies ······································································ 29

3.2.3.3. Thermodynamic parameters ···················································· 29

CHAPTER 4

RESULTS ·········································································· 31

4.1. BIOSORPTION STUDIES ON S. communis ·································· 32

4.1.1. BIOSORPTION CAPACITY OF CADMIUM BY

S. communis ··········································································· 31

4.1.1.1. Effect of pH ·········································································· 31

4.1.1.2. Effect of Contact Time ···························································· 31

4.1.1.3. Effect of Temperature ····························································· 31

4.1.1.4. Effect of Biosorbent Quantity ··················································· 32

4.1.1.5. Biosrption Equilibrium Isotherm ··············································· 32

4.1.1.5.1. Langmuir Isotherm ···································································· 32

4.1.1.5.2. Freundlich Isotherm ·································································· 33

4.1.1.5.3. Temkin isotherm ······································································ 33

4.1.1.6. Kinetic Studies ···································································· 34

4.1.1.7. Thermodynamics studies························································ 34

4.1.2. BIOSORPTION CAPACITY OF CHROMIUM BY

S. communis ··········································································· 35

4.1.2.1. Effect of pH ·········································································· 35

4.1.2.2. Effect of Contact Time ··························································· 35

4.1.2.3. Effect of Temperature ···························································· 35


4.1.2.5. Biosrption Equilibrium Isotherm ·············································· 36

4.1.2.5.1. Langmuir Isotherm ····································································· 36

4.1.2.5.2. Freundlich Isotherm ··································································· 36

4.1.2.5.3. Temkin isotherm ········································································ 37

4.1.2.6. Kinetic Studies ····································································· 37


4.2. BIOSORPTION STUDIES ON C. delmatica ········································ 38

4.2.1. BIOSORPTION CAPACITY OF CADMIUM BY C. delmatica ··············· 38

4.2.1.1. Effect of pH ·········································································· 38

4.2.1.2. Effect of Contact Time ···························································· 39

4.2.1.3. Effect of Temperature ····························································· 39

4.2.1.4. Effect of Biosorbent Quantity ···················································· 39

xiv

4.2.1.5. Biosrption Equilibrium Isotherm ··············································· 40



4.2.1.5.3. Temkin isotherm ········································································ 40

4.2.1.6. Kinetic Studies ····································································· 41


4.2.2. BIOSORPTION CAPACITY OF CHROMIUM BY C. delmatica ··············· 41

4.2.2.1. Effect of pH ·········································································· 41







4.2.2.5.3. Temkin isotherm ········································································ 43

4.2.2.6. Kinetic Studies ····································································· 43


4.3. BIOSORPTION STUDIES ON Spirogyra Spp. ···································· 44

4.3.1. BIOSORPTION CAPACITY OF CADMIUM BY

Spirogyra Spp. ········································································· 44

4.3.1.1. Effect of pH ·········································································· 44







4.3.1.5.3. Temkin isotherm ········································································ 46

4.3.1.6. Kinetic Studies ····································································· 47


4.3.2. BIOSORPTION CAPACITY OF CHROMIUM BY

Spirogyra Spp. ········································································· 47

4.3.2.1. Effect of pH ·········································································· 47






xv


4.3.2.5.3. Temkin isotherm ········································································ 49

4.3.2.6. Kinetic Studies ····································································· 49


4.4. STUDY OF BIOSORPTIVE BINDING SITES ON ALGAL SURFACE ······ 50

4.4.1. FTIR Spectroscopy ···································································· 50

4.4.1.1. Studies with S. communis ······················································· 50

4.4.1.2. Studies with C. delmatica ······················································· 51

4.4.1.3. Studies with Spirogyra Spp. ···················································· 51

4.4.2. Scanning Electron Microscopy (SEM) ··············································· 52

CHAPTER 5

DISCUSSION································································································································································ 134

5.1. BIOSORPTION CAPACITY OF THE

BIOSORBENTS ·········································································································································· 134

5.1.1. Effect of pH ······················································································ 134

5.1.2. Effect of Contact Time ········································································ 138

5.1.3. Effect of Temperature ········································································· 138

5.1.4. Biosorbent concentration ····································································· 139

5.1.5. Biosorption Equilibrium Isotherms ························································· 140

5.1.6. Kinetic Studies of Biosorption ······························································· 142

5.1.7. Thermodynamic studies ······································································ 144

5.2. BINDING SITES FOR BIOSORPTION

CAPACITY OF METAL IONS ········································································································ 145

5.2.1. FTIR Spectroscopy ············································································ 145

5.2.2. Scanning Electron Microscopy (SEM) ············································ 146

5.2.3. Biochemistry of Biosorptive Surface ······················································· 147

5.2.4. Mechanism of Biosorption ··································································· 147

Future Directions……………………………………………………………………..150

Conclusion…………………………………………………………………………….152

CHAPTER 6

REFERENCES ·························································································································································· 154

xvi

List of Tables 1. Table 3.1. Composition of Bold 1NV medium 54

2. Table 4.1. Langmuir Isotherm constants for the sorption of

Cd (II) and Cr (VI) ions by the investigated algal biosorbent. 55

3. Table 4.2. Freundlich Isotherm constants for the sorption

of Cd (II) and Cr (VI) ions by the investigated algal biosorbent. 56

4. Table 4.3. Temkin Isotherm constants for the sorption of

Cd (II) and Cr (VI) ions by the investigated algal biosorbent. 57

5. Table 4.4. pseudo-second order kinetic constants for the

biosorption of Cd (II) and Cr (VI) ions by the investigated

algal biosorbent. 58

6. Table 4.5. Thermodynamic parameters, Enthalpy and

Entropy for the biosorption of Cd (II) and Cr (VI)

ions by the investigated algal biosorbent. 59

7. Table 4.6. Thermodynamic parameters,

Arhenius constant (A°) and energy of activation (Ea)

for the biosorption of Cd (II) and Cr (VI) ions by the

investigated algal biosorbent. 60

8. Table 4.6. -∆G°(KJmol-1) values for biosorption of Cr (VI) and Cd (II)

ions by the investigated algal biosorbent at varying temperatures 61

9. Table 4.7 FTIR absorption wavelengths (Cm-1) for Cd (II) ion and Cr (IV) ions

biosorption by the investigated algal biomass 62

xvii

List of Graphs 1. Figure 4.1. Comparison of Biosorption of Cd (II) ion at varying pH in S. communis 63

2. Figure 4.2. Effect of contact time on biosorption of Cd (II) by S. communis 64

3. Figure 4.3. Effect of temperature on biosorption of Cd (II) ion by S. communis 65

4. Figure 4.4. Effect of biosorbent concentration on biosorption of Cd (II) ion by S.

communis 66

5. Figure 4.5. Langmuir isotherm for sorption of Cd (II) ion onto S. communis. 67

6. Figure 4.6. Freundlich isotherm for sorption of Cd (II) ion onto S. communis. 68

7. Figure 4.7. Temkin isotherm for sorption of Cd (II) ions onto S. communis. 69

8. Figure 4.8. Pseudo-second order kinetic graph for sorption of Cd (II) ions onto S.

communis. 70

9. Figure 4.9. Thermodynamic profile of sorption of Cd (II) ions onto S. communis for ∆H

and ∆S. 71

10. Figure 4.10. Thermodynamic profile of sorption of Cd (II) ions onto S. communis for Ea

and Ao. 72

11. Figure 4.11. Comparison of Biosorption of Cr (VI) ions at varying pH in S. communis 73

12. Figure 4.12. Effect of contact time on biosorption of Cr (VI) by S. communis 74

13. Figure 4.13. Effect of temperature on biosorption of Cr (VI) ions by S. communis 75

14. Figure 4.14. Effect of biosorbent concentration on biosorption of Cr (VI) ions by S.

communis 76

15. Figure 4.15. Langmuir isotherm for sorption of Cr (VI) ions onto S. communis. 77

xviii

16. Figure 4.16. Freundlich isotherm for sorption of Cr (VI) ions onto S. communis. 78

17. Figure 4.17. Temkin isotherm for sorption of Cr (VI) ions onto S. communis. 79

18. Figure 4.18. Pseudo-second order kinetic graph for sorption of Cr (VI) ions onto S.

communis. 80

19. Figure 4.19. Thermodynamic profile of sorption of Cr (VI) ions onto S. communis for ∆H

and ∆S. 81

20. Figure 4.20. Thermodynamic profile of sorption of Cr (VI) ions onto S. communis for Ea

and Ao. 82

21. Figure 4.21. Comparison of Biosorption of Cd (II) ions at varying pH in C. delmatica 83

22. Figure 4.22. Effect of contact time on biosorption of Cd (II) by C. delmatica 84

23. Figure 4.23. Effect of temperature on biosorption of Cd (II) ions by C. delmatica 85

24. Figure 4.24. Effect of biosorbent concentration on biosorption of Cd (II) ions by C.

delmatica 86

25. Figure 4.25. Langmuir isotherm for sorption of Cd (II) ions onto C. delmatica. 87

26. Figure 4.26. Freundlich isotherm for sorption of Cd (II) ions onto C. delmatica. 88

27. Figure 4.27. Temkin isotherm for sorption of Cd (II) ions onto C. delmatica 89

28. Figure 4.28. Pseudo-second order kinetic graph for sorption of Cd (II) ions onto C.

delmatica 90

29. Figure 4.29. Thermodynamic profile of sorption of Cd (II) ions onto C. delmatica for ∆H

and ∆S. 91

xix

30. Figure 4.30. Thermodynamic profile of sorption of Cd (II) ions onto C. delmatica for Ea

and Ao. 92

31. Figure 4.31. Comparison of Biosorption of Cr (VI) ions at varying pH in C. delmatica

93

32. Figure 4.32. Effect of contact time on biosorption of Cr (VI) by C. delmatica 94

33. Figure 4.33. Effect of temperature on biosorption of Cr (VI) ions by C. delmatica 95

34. Figure 4.34. Effect of biosorbent concentration on biosorption of Cr (VI) ions by C.

delmatica 96

35. Figure 4.35. Langmuir isotherm for sorption of Cr (VI) ions onto C. delmatica 97

36. Figure 4.36. Freundlich isotherm for sorption of Cr (VI) ions onto C. delmatica 98

37. Figure 4.37. Temkin isotherm for sorption of Cr (VI) ions onto C. delmatica 99

38. Figure 4.38. Pseudo-second order kinetic graph for sorption of Cr (VI) ions onto C.

delmatica 100

39. Figure 4.39. Thermodynamic profile of sorption of Cr (VI) ions onto C. delmatica for ∆H

and ∆S. 101

40. Figure 4.40. Thermodynamic profile of sorption of Cr (VI) ions onto C. delmatica for Ea

and Ao. 102

41. Figure 4.41. Comparison of Biosorption of Cd (II) ion at varying pH in Spirogyra Spp. 103

42. Figure 4.42. Effect of contact time on biosorption of Cd (II) by Spirogyra Spp. 104

43. Figure 4.43. Effect of temperature on biosorption of Cd (II) ion by Spirogyra Spp. 105

44. Figure 4.44. Effect of biosorbent concentration on biosorption of Cd (II) ion by Spirogyra

Spp. 106

45. Figure 4.45. Langmuir isotherm for sorption of Cd (II) ion onto Spirogyra Spp. 107

xx

46. Figure 4.46. Freundlich isotherm for sorption of Cd (II) ion onto Spirogyra Spp. 108

47. Figure 4.47. Temkin isotherm for sorption of Cd (II) ions onto Spirogyra Spp. 109

48. Figure 4.48. Pseudo-second order kinetic graph for sorption of Cd (II) ions onto

Spirogyra Spp. 110

49. Figure 4.49. Thermodynamic profile of sorption of Cd (II) ions onto Spirogyra Spp. for

∆H and ∆S. 111

50. Figure 4.50. Thermodynamic profile of sorption of Cd (II) ions onto Spirogyra Spp. for Ea

and Ao. 112

51. Figure 4.51. Comparison of Biosorption of Cr (VI) ions at varying pH in Spirogyra Spp.

113

52. Figure 4.52. Effect of contact time on biosorption of Cr (VI) by Spirogyra Spp. 114

53. Figure 4.53. Effect of temperature on biosorption of Cr (VI) ions by Spirogyra Spp. 115

54. Figure 4.54. Effect of biosorbent concentration on biosorption of Cr (VI) ions by

Spirogyra Spp. 116

55. Figure 4.55. Langmuir isotherm for sorption of Cr (VI) ions onto Spirogyra Spp. 117

56. Figure 4.56. Freundlich isotherm for sorption of Cr (VI) ions onto Spirogyra Spp. 118

57. Figure 4.57. Temkin isotherm for sorption of Cr (VI) ions onto Spirogyra Spp. 119

58. Figure 4.58. Pseudo-second order kinetic graph for sorption of Cr (VI) ions onto

Spirogyra Spp. 120

59. Figure 4.59. Thermodynamic profile of sorption of Cr (VI) ions onto Spirogyra Spp. for

∆H and ∆S. 121

xxi

60. Figure 4.60. Thermodynamic profile of sorption of Cr (VI) ions onto Spirogyra Spp. for

Ea and Ao. 122

61. Figure 4.61. (a) SEM of S. communis (Control) 123

62. Figure 4.61. (b) SEM of S. communis (Cd-loaded) 123

63. Figure 4.61. (a) SEM of S. communis (Cr- loaded) 123

64. Figure 4.62. (a) SEM of C. delmatica (Control) 124

65. Figure 4.62. (b) SEM of C. delmatica (Cd-loaded) 124

66. Figure 4.62. (a) SEM of C. delmatica (Cr- loaded) 124

67. Figure 4.63. (a) SEM of Spirogyra Spp. (Control) 125

68. Figure 4.63. (b) SEM of Spirogyra Spp. (Cd-loaded) 125

69. Figure 4.63. (a) SEM of Spirogyra Spp. (Cr- loaded) 125

70. Figure 4.64. (a) FTIR Spectra of S. communis (Control) 126

71. Figure 4.64. (b) FTIR Spectra of S. communis (Cd-loaded) 126

72. Figure 4.64. (c) FTIR Spectra of S. communis (Cr-loaded) 126

73. Figure 4.65. (a) FTIR Spectra of C. delmatica (Control) 127

74. Figure 4.65. (b) FTIR Spectra of C. delmatica (Cd-loaded) 127

75. Figure 4.65. (c) FTIR Spectra of C. delmatica (Cr-loaded) 127

76. Figure 4.66. (a) FTIR Spectra of Spirogyra Spp. (Control) 128

77. Figure 4.66. (b) FTIR Spectra of Spirogyra Spp. (Cd-loaded) 128

78. Figure 4.66. (c) FTIR Spectra of Spirogyra Spp. (Cr-loaded) 128

79. Figure 5.1. Involvement of functional groups in providing divalent metal ions

(M2+) a site of attachment 129

80. Figure 5.2. Incorporation of Cr (VI) ion in the organic components on the cell wall

of green algae 130

81. Figure 5.3. The carboxyl ion on the surface of algae provide site of biosorption onto the surface

of biosorbets, studied in the present investigation 131

82. Figure 5.4. The mechanism of ion exchange at electonegative functional group 132

83. Figure 5.5. The fibrillar molecules of algal cell walls of green algae 133

1

CHAPTER 1

INTRODUCTION

―The quality of life on earth is inextricably linked to the overall quality of

the environment. On the rise is the understanding that there has been

negligence and carelessness in human activities concerning some negative

impacts of rapid industrialization as well as exploration and utilization of some

natural resources, especially during the past few decades. The indiscriminate

release of hazardous pollutants by industries and dumping of domestic sewage

in the human environment pose a major threat to all kinds of organisms

inhabiting aquatic as well as terrestrial ecosystems (Filippis and Pallaghy,

1994). In this context, a variety of industrial effluents are routinely being

discharged into the environment, including heavy metal pollutants, such as

cadmium (Cd), chromium (Cr), lead (Pb), mercury (Hg), etc. This problem is

worldwide, and the estimated number of contaminated sites with heavy metals

is increasing constantly in significant numbers (Cairney, 1993).‖

―Generally, metals that have a density of 5 g/cm3 or greater are

considered as heavy metals. In the periodic table, among the transition

elements from group V (excluding Sc and Ti) to the half-metal As (termed as a

metalloid), from Zr (but not Y) to Sb, from La to Po, the Lanthanides and the

Actinides can be referred to as "heavy metals". Of the 90 naturally occurring

elements, 21 are non-metals, 16 are light metals, and the remaining 53 (with As

included) are heavy metals (Weast, 1984). Most heavy metals are transition

elements with incompletely filled d-orbitals. These d-orbitals provide heavy

metal cations with the ability to form complex compounds some of which may

be redox active (Nies, 1999). In general, living system requires some of these

elements (metal ions) for metabolic activities but only in trace amounts. The

intake of these metals in greater concentration than required by the organisms

may cause health hazards. It is now established that one of the most

dangerous and pernicious forms of pollution arises from the potential

mobilization of a spectrum of toxic trace metals, such as Cr, Cd, Pb, Hg, etc.,

and metalloids (e.g., As) in our environment (Volesky, 2007). For example, both

Cr and Cd accumulate in the environment in large concentrations and are

2

increasingly becoming a threat to the ecosystem in addition to directly affecting

human health (U.S. EPA, 1997). ―

―Cadmium is a soft silver-white metal that is usually found in combination

with other elements, its compounds range in solubility in water from quite

soluble to practically insoluble (Avery et al., 1998) and its atomic weight is

112.41 g/mol (ATSDR,1997). This heavy metal is used to manufacture

pigments and batteries and in the metal-plating and plastics industries (ATSDR,

1997). The U.S. Environmental Protection Agency (U.S.EPA, 1998) has

classified Cd as a Group B1, probable human carcinogen. The largest sources

of airborne Cd in the environment are the burning of fossil fuels, such as coal or

oil, and incineration of municipal waste materials. Cadmium may also be

emitted into the air from zinc, lead, or copper smelters (ATSDR, 1997). For

nonsmokers, food is generally the largest source of Cd exposure. Cadmium

levels in some foods can increase by the application of phosphate fertilizers or

sewage sludge to farm fields (ATSDR, 1997). Smoking is another important

source of human exposure to Cd. Smokers have about twice as much Cd in

their bodies as do nonsmokers (ATSDR, 1997). In animals, acute inhalation

exposure to high levels of Cd in humans may result in adverse effects on the

lungs, such as bronchial and pulmonary irritation. Acute exposure to high level

of Cd can result in long-lasting impairment of lung function (ATSDR, 1997;

U.S.EPA, 1999). Cadmium is considered to have high acute toxicity, based on

short-term animal tests, such as the LC50 test in rats (U.S.EPA, 1999). Chronic

inhalation and oral exposure of humans to Cd (II) results in a build-up of Cd (II)

level in the kidneys that can cause kidney disease, including proteinuria, a

decrease in glomerular filtration rate, and an increased frequency of kidney

stone formation (ATSDR, 1997; U.S.EPA, 1999). In humans, chronic exposure

to Cd can affect lungs, causing bronchiolitis and emphysema (ATSDR, 1997;

U.S.EPA, 1999). In additon, it also affects liver, bone, immune system, blood,

and nervous system adversely (ATSDR, 1997; U.S. EPA, 1999) . The

Reference Dose (RfD) for Cd in drinking water is 0.0005 mg/kg/d, and the RfD

for dietary exposure to Cd is 0.001 mg/kg/d; both are based on significant

proteinuria in humans. The RfD is an estimate (with uncertainty spanning

perhaps an order of magnitude) of a daily oral exposure to the human

population (including sensitive subgroups) that is likely to be without

3

appreciable risk of deleterious non-cancer effects during a lifetime. It is not a

direct estimator of risk, but rather a reference point to gauge the potential

effects. At exposures greater than the RfD, the potential for adverse health

effects increases (U.S.EPA, 1999). The United States Environmental Protection

Agency (U.S.EPA) has high confidence in both RfD values based primarily on a

strong database for Cd toxicity in humans and animals that also permits

calculation of pharmacokinetic parameters of Cd absorption, distribution,

metabolism, and elimination (U.S.EPA, 1999). EPA has not established a

reference concentration (RfC) for Cd (U.S.EPA, 1999); however, the California

Environmental Protection Agency (Cal EPA) has established a chronic

reference exposure level of 0.00001 mg/m3 for Cd based on kidney and

respiratory effects in humans. The Cal EPA reference exposure level is a

concentration at or below which adverse health effects are not likely to occur

(U.S. EPA, 1999). Animal studies provide evidence that Cd has developmental

effects, such as low fetal weight, skeletal malformations, interference with fetal

metabolism, and impaired neurological development, via inhalation and oral

exposure (ATSDR, 1997; U.S.EPA, 1999). Decreased reproduction and

testicular damage have been noted following oral exposures (ATSDR, 1997).

Several occupational studies have reported an excess risk of lung cancer in

humans from exposure to inhaled Cd; however, the evidence is limited rather

than conclusive due to confounding factors. (ATSDR,1997; U.S. EPA,1999).

Animal studies have reported cancer resulting from inhalation exposure to

several forms of Cd, while animal ingestion studies have not demonstrated

cancer resulting from this heavy metal (ATSDR, 1997; U.S. EPA,1999). ―

―U.S. EPA estimates that if an individual were to continuously breathe air

containing Cd at an average of 0.0006 μg/m3 (6 x 10-7 mg/m3) over his or her

entire lifetime, that person would theoretically have no more than a one-in-a-

million increased chance of developing cancer as a direct result of breathing air

containing this heavy metal. Similarly, continuously breathing of air containing

0.006 μg/m3 (6 x 10-6 mg/m3) of Cd would result in not greater than a one-in-a-

hundred thousand increased chance of developing cancer, and air containing

0.06 μg/m3 (6 x 10-5 mg/m3) of this heavy metal would result in not greater than

a one-in-ten thousand increased chance of developing cancer (U.S. EPA,

1999). Earthworms and other essential soil organisms are extremely

4

susceptible to Cd poisoning. They can die at very low concentrations and this

has consequences for soil biota and soil structure. When Cd concentrations in

soils are high they can influence soil processes of microorganisms and threat

the whole soil ecosystem. In aquatic ecosystems, Cd can bio-accumulate in

mussels, oysters, shrimps, lobsters and fish. The susceptibility to Cd can vary

greatly between aquatic organisms. Salt-water organisms are known to be

more resistant to Cd poisoning than freshwater organisms (Naja and Volesky,

2009).‖

―The other heavy metal, chromium (Cr), is a steel-gray solid with a high

melting point (1907oC) and an atomic weight of 51.996 g/mol. Chromium has

oxidation states ranging from chromium [Cr (VI)] to hexavalent [Cr (VI)] form of

Cr (ATSDR, 1998). Chromium forms a large number of compounds, in both the

trivalent Cr3+ and the hexavalent chromium [Cr (VI)] forms. Chromium

compounds are stable in the trivalent state, with the hexavalent form being the

second most stable state (ATSDR, 1998). The trivalent chromium (Cr3+)

compounds are sparingly soluble in water and may be found in water bodies as

soluble Cr (III) complexes, while the Cr (VI) compounds are readily soluble in

water (ATSDR, 1998). A number of studies in humans and animals have been

reported to show the effect of Cr (VI). For example, in humans exposure to Cr

(VI) via inhalation may result in complications during pregnancy and childbirth

(ATSDR, 1998). Oral studies have reported severe developmental effects in

mice, such as gross abnormalities and reproductive effects including decreased

litter size, reduced sperm count, and degeneration of the outer cellular layer of

the seminiferous tubules. (ATSDR, 1998; U.S.EPA, 1998b). Epidemiological

studies of workers have clearly established that inhaled chromium is a human

carcinogen, resulting in an increased risk of lung cancer. Although chromium-

exposed workers were exposed to both Cr3+ and Cr (VI) compounds, only Cr

(VI) has been found to be carcinogenic in animal studies, thus EPA has

concluded that only Cr (VI) should be classified as a human carcinogen

(ATSDR, 1998; U.S. EPA, 1999). Animal studies have shown Cr (VI) to cause

lung tumors via inhalation exposure (ATSDR, 1998; WHO, 1987). EPA has

classified Cr (VI) as a Group A, known human carcinogen by the inhalation

route of exposure having estimated inhalation unit risk of 1.2 × 10-2 (µg/m3)-1

(U.S. EPA, 1999). If an individual were to continuously breathe air containing

5

Cr at an average of 0.00008 µg/m3 (8 x 10-8 mg/m3) over his or her entire

lifetime, that person would theoretically have no more than a one-in-a-million

increased risk of developing cancer(U.S. EPA, 1999). Similarly, the U.S. EPA

estimates that continuously breathing air containing 0.0008 µg/m3 (8 x 10-7

mg/m3) would result in not greater than a one-in-a-hundred thousand increased

risk of developing cancer during one's lifetime, and air containing 0.008 µg/m3

(8 x 10-6 mg/m3) of Cr would result in not greater than a one-in-ten-thousand

increased risk of developing cancer during one's lifetime (U.S. EPA, 1999).

Recent studies suggest that Cr (VI) is responsible for most of the toxic actions,

although much of the underlying molecular damage may be due to its

intracellular reduction to the even more highly reactive and short-lived chemical

forms, Cr3+ and Cr2+ (Leusch et al., 1995; Seiki and Suzuki, 1998). Exposure to

Cr (VI) can result in various point mutations in DNA, chromosomal damage,

and oxidative changes in proteins (Dayan, 2001). Biochemical studies of the

DNA-damaging effects and of the pathogenesis of the allergic reactions to Cr

are still not completely understood. ‖

―These metal pollutants, contained in industrial effluents and domestic

sewage, being dumped into the water bodies (streams, rivers, oceans, etc.),

tend to accumulate in the biomass of aquatic flora and fauna. In addition, crops

and other terrestrial plants also get contaminated mainly through irrigation.

These toxic metals tend to budget up as they travel through the food chain,

posing a serious threat to the animals and humans. Thus, the entire

environment is at stake (Ellwood et al., 1989; Leusch et al., 1995; Patra et al.,

2004). ‖

―Thus, the present scenario presents an urgent need for necessary

action towards the removal of toxic heavy metals including Cr (VI) and Cd (II) at

scientific platform. Recently, necessary emphasis has been placed to tackle

this alarming situation and a number of studies have been reported in this

connection.

―Adsorption process has emerged as the most promising technique and

has attracted the attention of many researchers in this field. Both natural

inorganic and organic adsorbent materials have been tested in this regard. For

example, different natural oxides were used in heavy metal removal, such as

aluminum oxide in Cr (VI), Cd (II), and Pb2+ removal from wastewaters, geothite

6

or iron oxide coated with sand and clay for Cd (II), and Pb2+ removal, and

removal of Pb2+ compounds using ZnO loaded onto activated carbon (Gupta

and Tiwari, 1985; Srivastava et al., 1989; Bailey et al., 1992; Potgieter et al.,

2006; Kikuchi et al., 2006). On the other hand, a great variety of synthetic

adsorbents, such as Silica gel and TiO2 were tested for heavy metal removal

(Chakravarty et al., 2002; Ahluwalia and Goyal 2007; Bailey et al., 1999;

Varshney et al., 2003; Jiang et al., 2007). Silica gel is a synthetic amorphous

polymer that has silanol groups on its surface that allow metal adsorption

(Ricordel et al., 2001; Michard et al., 1994). Moreover, the modification of the

silica gel surface can be used to enhance its adsorption properties (Tran et al.,

1999; Kocjan et al., 1996; Bresson et al., 1998). On the other hand, TiO2-

mediated photocatalytic treatment of metal-chelate waste has also been

reported (Chiron et al., 2003; Prairie et al., 1993a, 1993b; Madden et al., 1997).

The oxidation process occurs either at the TiO2 surface or at a small distance

from the surface in the solution phase (Davis and Green, 1999). TiO2-SiO2

mixed oxide has also been tested successfully as sorbent for removal of heavy

metals to get the benefits of both oxides (i.e., the high surface area of silica as

well as the photo-catalytic ability of TiO2) (Ismail et al., 2005). ‖

―Although many of natural and/or synthetic adsorbents can effectively

remove dissolved heavy metals, most of them show some disadvantages, such

as poor adsorption capacity, a low efficiency/cost ratio, and ineffectiveness for

low metal concentrations (Ismail et al., 2005). Commonly used synthetic

adsorbents are based on various physicochemical mechanisms, such as

chemical precipitation, oxidation/reduction, evaporation, reverse osmosis, etc.,

but most of them have some demerits in their practical application. For

example, chemical precipitation is not effective in removing heavy metals

especially in low concentrations; chemical oxidation and reduction mechanisms

are slow in action and climate sensitive; electrochemical treatment, evaporation

and reverse osmosis are expensive processes; adsorption is not effective in

case of metal ions (Volesky et al., 2000). In addition, the price of ion exchange

resins, that are hydrocarbon derivatives, is invariably linked to that of crude oil.

Needless to say, crude oil is a finite resource and, in addition to that

disadvantage, its price is also very much subject to the world trading stability.

This uncertainty guarantees the possibility of easily opening search for

7

alternatives (such as biosorption), which would be better and effective as well

(Volesky, 2007). In this context, use of organisms to remove toxic metals from

wastewater (bioremediation) has been advocated. For example, various types

of organisms including algae (green algae, brown algae, and red algae),

bacteria, yeast and fungi (such as Rhizopus oligoorus, Datura innoxia,

Sacharomycete cerevisiae) have been studied for their capacity to remove

heavy metal ions from solution (Ahalya et al., 2005; Chong and Volesky,1990;

Drake et al., 1996; Godlewska-Zylkiewicz, 2006; Tobin et al., 1984; Volesky,

1990, 2001, 2007; Leusch et al., 1995; Figueira et al., 1997; Matheickal et

al.,1999; Ozer et al., 2000; Lee et al., 2000; Cervantes et al., 2001; Terry and

Stone, 2002; Akhtar et al., 2003; Davis et al., 2003a,b; Andrade et al., 2005).

Among the various organisms, algae stand out as the most promising

candidate for the removal of heavy metals (Chong and Volesky, 1996; Leusch

et al., 1995; Lee et al., 2000; Cervantes et al., 2001; Davis et al., 2003a,b;

Arica and Bayramoglu, 2005; Chojnacka et al.,2004; Mehta and Gaur, 2005;

Vilar et al., 2005; Volesky, 1990, 2000, 2007). ‖

―The application of biosorption technology has emerged as a significant

research area in the past decade. Biosorption, here refers to the use of dead

biomass for the removal of heavy metal pollutants. It must be distinguished

from metabolically driven active uptake of metals by living organisms, the

phenomenon called bioaccumulation (Seki and Suzuki, 1998). The

accumulation of heavy metals in algae involves two processes: an initial rapid

(passive) uptake followed by a much slower (active) uptake (Bates et al., 1982;

Arief et al., 2008). During the passive uptake, metal ions adsorb onto the cell

surface within a relatively short span of time (few seconds or minutes), and the

process is metabolism independent. Active uptake is metabolism-dependent,

causing the transport of metal ions across the cell membrane into the

cytoplasm. In some instances, the transport of metal ions may also occur

through passive diffusion owing to metal-induced increase in permeability of the

cell membrane (Gadd, 1988). ‖

―The process of accumulation and adsorption of metals by algae

involves adsorption onto the cell surface (wall, membrane or external

polysaccharides) and binding to cytoplasmic ligands, phytochelatins and

metallothioneins, and other intracellular molecules. Localization of metal ions

8

on algal cell has been carried out by electron microscopy and X-ray energy

dispersive analysis studies (Mehta and Gaur, 2005). Spectroscopy has been

used for determining the oxidation state of bound metal on algal cell.

Transmission electron microscopy (TEM) has shown cell wall as the most likely

location of Cd adsorption by Ectocarpus siliculosus (Winter, et al., 1994).

Scanning electron microscopy (SEM), in combination with X-ray microanalysis,

has clearly revealed that most of the sites for metal sorption are present on the

surface of algal cells (Klimmek et al., 2001). The algal cell wall has many

functional groups, such as hydroxyl (OH), phosphoryl (PO3O2), amino (NH2),

carboxyl (COOH), sulphydryl (SH), etc., which confer negative charge to the

cell surface. Since metal ions in water are generally in the cationic form, they

are adsorbed onto the cell surface (Crist et al., 1981; 1991; Xue et al., 1988;

Romero-onzalez, et al., 2001; Skowron et al., 2002). Each functional group has

a specific pKa (dissociation constant) (Eccles, 1999; Niu and Volesky, 2001),

and it dissociates into corresponding anion and proton at a specific pH. These

functional groups are found associated with various cell wall components, e.g.,

peptidoglycan, teichouronic acid, teichoic acids, polysaccharides and proteins.

Because distribution and abundance of cell wall components vary among

different algal groups, the number and kinds of functional group involved in the

biochemistry of cell wall also varies in different algal groups. Among different

cell wall constituents, polysaccharides and proteins have most of the metal

binding sites (Kuyucak and Volesky, 1989a, Lawal, et al., 2010). ‖

―The cell wall of various groups of algae has been evaluated for their

capacity to bind and remove heavy metals from the environment. For example,

the cell wall of green algae contains heteropolysaccharides, which offer

carboxyl and sulfate groups (Lee, 1980) for sequestration of heavy metal ions.

The extracted polysaccharides (12% of dry weight) from Ulva sp. contained

16% sulfate and 15–19% uronic acid (McKinnel and Percival, 1962). Protein

content of cell wall of green algae ranges from 10 to 70% (Siegel and Siegel,

1973). Aspartic and glutamic acid account for 12% of protein content of cell

wall, which correspond to 0.15 meq g−1 of carboxylic groups per dry weight. In

the cell wall of green algae, lysine and arginine make up 13% of protein (0.08

meq g−1 of amino groups) (Chapman, 1980; Lee, 1980; Guing and Blunden,

1991). Marine algae have been the focus of numerous biosorption studies and

9

their excellent metal binding capacity has been widely acknowledged. The main

constituents of the cell wall of brown algae are cellulose, as the fibrous

skeleton, and alginate and fucoidan that constitute the amorphous matrix, and

extracellular mucilage (Lee, 1980). It is fairly well recognized that alginate is

mainly involved in metal accumulation by brown algae (Kuyucak and Volesky,

1989b). Alginate is defined as the ammonium or alkali salt of alginic acid.

Following a mild acid hydrolysis of alginic acid, three kinds of segments are

demonstrated as building blocks of the polymer. These are D-mannuronic acid

and L-guluronic acid units, and alternating D-mannuronic acid and L-guluronic

acid residues (Davis et al., 2003a, b). The carboxyl groups of each polymer

segment may play important role as the site for cation binding. Polyguluronic

acid shows a high specificity for divalent metal ions (Puranik, et al., 1999). The

affinity of alginates for divalent cations, such as Pb (II), Cu (II), Cd (II), Zn (II),

Ca (II), etc., increased with guluronic acid content. Alginate content of brown

algae is 10 to 40% of dry weight (Percival and McDowell, 1967). Alginate

concentration in Sargassum fluitans is 45% of its dry weight, corresponding to

2.25 mmol of carboxyl groups g−1 of biomass (Fourest and Volesky, 1996).

Calpomenia has 5–14% alginate (0.25–0.7 meq g−1 carboxyl groups)

(Kalimuthu et al., 1991). Involvement of carboxyl groups in adsorption of

cations, such as Cu onto Chlorella vulgaris (Seki et al., 2000) suggested that

sorption of bivalent metal ions by a marine microalga, Heterosigma akashiwo,

involves mono-dentate binding on carboxylic and phosphotidic type sites

(Mehta et al., 2002). A significant role of carboxyl groups in sorption of heavy

metals has also been very well demonstrated in fungi as well as in higher

plants (Kapoor and Viraraghavan, 1997; Aksu and Donmez, 2001; Lawal et al.,

2010). This declares algae as an efficient potential biosorbents. Certainly, over

the years, use of dead algae instead of live algae has gained popularity in the

process of biosorption for the removal of heavy metal ions from solutions. Here,

biomass is regarded as assemblage of various polymers, such as cellulose,

pectin, glycoproteins, etc., capable of binding to toxic heavy metal cations

declaring them strong candidate adsorbents that could possibly be used in

wastewater treatment for the removal of toxic heavy metal ions, hence, opening

the doors of providing a better and cost-effective alternatives solutions

(Volesky, 2007, Arief et al., 2008). ‖

10

―The potential of filamentous green algae (Chlorophyceae) in removing

metal ions from solutions remains largely unexplored, except for a few reports

(Aksu et al., 1998; Ozer et al., 2000; Nuhoglu et al., 2002; Mohapatra and

Gupta, 2005). Now, they are being predicted as leading candidates in this race

due to the biochemistry of their cell wall (Tereshina et al., 1999; Macfie and

Welbourn, 2000; Hung et al., 2003; Andrade et al., 2005; Arica and

Bayramoglu, 2005; Arief et al., 2008; Pahalvanzadeh et al., 2010). Freshly

available data regarding the capacity of filamentous green algae, such as

Spirogyra and Cladophora to adsorb heavy metal ions, such as Cd and Cu, etc.

has opened the doors to new optimism (Singh et al., 2006). They (filamentous

green algae including Spirogyra and Cladophora) are easily available in terms

of amount of biomass especially in tropical freshwater bodies. These algae are

relatively newly known in this context and their potential to establish as better

and cost-effective biosorbents in this scenario needs further exploration. ‖

―As model predictions of heavy metal biosorption become more

sophisticated, there is an underlying need to appreciate the basic cell biology

and biochemistry of the filamentous green algae including the genera Spirogyra

and Cladophora and their comparison with other algae. Spectroscopic studies

on chemistry of cell wall need to be focused to enhance the utilization of

biosorption technology due to the capacity of cell wall of electrostatic attraction

and complexation with metal cation (Lee, 1980). Algal cell walls are comprised

of a fibril skeleton and an amorphous embedding matrix. The most common

fibrillar skeleton material is cellulose (Davis et al., 2003b). The well-known

understanding of filamentous green algae regarding their ability to bind metal

ions and the functional groups in the cellulose of their cell wall makes them

potential candidates as priority biosorbents for toxic metal cations. ‖

―The present study deals with the biosorption of heavy metal ions, such

as Cd and hexavalent chromium (Cr (VI)) by using the biomass of green algae:

Spirogyra and Cladophora. Research at cellular level to improve the

understanding of possible binding sites in the cell wall of the used biosorbents

was undertaken to understand the extent of biosorption and to determine as to

how various polysaccharides and peptides are possibly involved in binding with

Cd (II) and Cr (VI).

11

CHAPTER 2 REVIEW OF LITERATURE

2.1 BASIC UNDERSTANDING

―Understanding of contamination of environment by heavy metal

pollutants and its remediation by using various synthetic and biological

adsorbent is largely under investigation by the scientific community. As early as

in 1983, Ramaswamy et al., on the basis of some studies pointed out that algae

are probable candidate for removing heavy metal pollutants from the

environment, since then scientists have not only evaluated pollution status of

aquatic ecosystems contaminated with heavy metals but also numerous

remedial studies have been carried out by using algae to mitigate the heavy

metal pollutants from the environment.‖

―It is estimated that over a billion human beings are exposed to rather

elevated concentrations of toxic metals, such as Cadmium (Cd), Chromium

(Cr), Copper (Cu), Lead (Pb) and Zinc (Zn) etc., and metalloids, such as

Arsenic(As) in the environment, and many of them are suffering from sub-

clinical metal poisoning (Nriagu, 1980). Ramaswamy et al., (1984) studied that

electroplating industrial wastes contained heavy metal toxins, such as Cd, Cr,

Cu, Pb, and Zn, in objectionable amounts and also reported accumulation of

heavy metals in some algal taxa. Jordan (1993), using atomic absorption

spectrophotometer technique, evaluated the contamination from tannery

discharges into rivers in the mining state of Gerais, Brazil in sediments, water

and fish samples collected from the rivers and reported concentration of Cr in

the biotic and abiotic components of ecosystem; fish possessed approximately

35 times more heavy metal content than it should have under normal

conditions. Leslie et al., (2001) studied the impact of Cr contamination on the

benthic macro-invertebrate community of the Chusovaya River in the Ural

Mountains of Russia. Chemical analysis of water, sediments and detritus

indicated that the main contaminant present was indeed Cr and that the levels

of Cr concentrations were about 450 times higher in water and 25 times higher

in the sediments compared with a clean reference site upstream. The results

indicated that the observed extreme Cr contamination had an adverse effect on

12

aquatic life in the Chusovaya River, both at the community level (reduced

diversity) as well as at the level of individuals (sublethal effects on surviving

individuals). Szymanowska, et al., (1999) measured concentration of heavy

metals (Ni, Cr, Co, Zn, Mn, Pb, Cd, Cu, Hg, Fe, etc.) as well as macronutrients

in water, bottom sediments and plants of three lakes in West Poland. These

plants contained elevated levels of Zn, Cd, Cr, and Hg. Analysis of water and

bottom sediments indicated that the lakes were polluted with Zn, Cd, Cu, and

Pb and partly with Ni, and Hg. In India, Reddy (1995) investigated paper mill

effluents discharged into a river and their effects on algae and reported that

blue green algae seemed to be more tolerant than green algae and the

concentration of various heavy metals analyzed was in order of

Zn>Cu>Pb>Ni>Co>Mn. Burdin and Bird (1994) studied the heavy metals

accumulation by carrageenan and agar producing algae; the heavy metals

measured in living and lyophilized algal thalli were Cu, Cd, Ni, Zn, Mn and Pb

and the agar producing macro algae were Gracilaira tikvahiae and Gelidium

pusillum. Garnham, et al., (1992) observed the accumulation of heavy metals

by microalgae and reported two phases of metal ion uptake; accumulation of

heavy metals by micro algae may consist of two phases: metabolism–

independent binding to cell walls/extracellular polysaccharides (biosorption)

followed or accompanied by intracellular uptake which may be energy

dependent. Chlorella salina was found to exhibit both uptake mechanisms for

Co, Mn, Zn and Calcium. Malea et al., (1994) reported bioaccumulation of

elevated levels of heavy, such as Fe, Cu, Zn, Cd, Pb, Na, K, Ca and Mg in

seven species of red algae (Rhodophyta) from Antikyra Gulf, Greece. Okamura

and Aoyama (1994) studied the relationship between an interactive toxic effect

and distribution of heavy metals in algal cells by culturing the green alga,

Chlorella ellipsoidea for 6 days in the presence of Cd or Cr and determined the

amounts of the Cd (II) ions accumulations in soluble fraction and in membrane

fraction as 50 and 20% respectively; also presence of Cr (VI) ions changed the

Cd (II) ions in both the fractions up to 40%. Ali et al., 1996 analyzed various

industrial effluents and found heavy metals and trace metal contents including

Cr, Cd, Cu, Pb, Ni, Zn, Co, Mg, Fe, and As in these effluents. Sial et al (2006)

declared that over a thousand industrial units (textile, pharmaceutical,

chemicals, food industries, ceramics, steel, oil mills, and leather tanning) in

13

Pakistan are responsible for dumping heavy metal pollutants to the natural

water bodies. These authors also reported that due to the lack of strict

environmental planning guidelines in Pakistan, industries and factories dump

their solid and liquid wastes in environments adjacent to these sites, sewers,

nallahs, and streams, these pollutants mix with groundwater thereby raising

levels of heavy metal pollutants much higher than the levels recommended by

WHO and also much higher than the permissible limits of NEQS (Gulfraz et al.,

1997; 2002). ‖

2.2. HAZARDS OF HEAVY METAL POLLUTANTS

―The alarming influx of heavy metals in the environment is a serious

threat to the living organisms including humans. Degraeve (1981) found that

exposure of heavy metals at higher concentrations induced tumors and

mutations in animals. Wagner (1993) reported the capacity of heavy metals to

cause genetic damage to germ cells of male and female animals including

humans. Groten and Van Bladeren (1994) declared heavy metal toxins as

cumulative toxins, which through biomagnification in plants could affect human

health. Irshad and Ali(1997) suggested that industrial pollutants, such as

organic and inorganic chemicals and toxic metals were increasingly polluting

the air, soil and water in Pakistan. Gabbrielli et al., (1990) reported that higher

concentrations of heavy metals in soils inhibit plant growth, nutrient uptake and

physiological and metabolic processes. Sanita and Gabbrielli (1999) stated that

heavy metal pollutants would results in chlorosis, damage to root tips, reduced

water and nutrient uptake and damage to enzyme system. Lannelli et al.,

(2002) stated that heavy metals, like other environmental stressors, also

induced increased antioxidant enzyme activities in plants. Shanker et al.,

(2003) reported metabolic alterations by Cr exposure in plants effecting the

enzymes and other metabolites and ability of Cr to generate reactive oxygen

species which may cause oxidative stress. Ahammadsahib et al., (1989),

reported the effect of Cd toxicity in animal body by affecting the enzyme system

and resulting in adverse effects on kidneys and lungs. Khan (2006) reported

that human exposure to toxic heavy metal ions through ingestion of

contaminated food or uptake of drinking water could lead to the accumulation of

these pollutants in humans, animals, and plants. ‖

14

2.3. REMOVAL OF TOXIC HEAVY METAL

―The present scenario of the pollution status regarding accumulation and

biomagnification of heavy metals, such as, Cr, Cd, Cu, Pb, Hg, Ni, etc. creates

an urgent need to mitigate these toxins from the environment. A number of

studies have been undertaken to resolve the matter. For example, various

methods based on some physico-chemical mechanisms by the use of synthetic

adsorbents have been popular over the years for the removal of these heavy

metals from wastewater. .Veglio and Beolchini (1997) reported an account of

various studies regarding the removal of heavy metal ions from waste solutions

and the recovery of precious metals by using various methodologies.

Ramaswamy et al., (1983) suggested the use of algal biomass to mitigate

heavy metal content and toxicity in the aquatic ecosystem. Saha et al., (2002)

studied the adsorption of toxic metal ions, such as Cd, Zn, and Pb by using

hydroxyaluminum- and hydroxyaluminosilicate-montmorillonite complexes.

Feng and Aldrich (2004) undertook a toxicological approach to determine the

heavy metal binding capacity of soils. Hammaini et al., (2006) reported the

biosorption and desorption studies of heavy metal including Cu2+, Cd, Zn, Ni

and Pb by using activated sludge by its chemical modification. ‖

2.4. BIOSORPTION AND BIOREMEDIATION

―In addition to synthetic biosorbents, the biomass of various

microorganisms and macro-organisms including algae, fungi, and bacteria has

emerged as a powerful tool for the removal of toxic metal ions from

wastewater. Avery et al., (1994) proposed the use of photosynthetic

microorganisms (micro algae) in the wastewater of industrial effluents as bio-

remedial measures. Aloysius et al., (1999) suggested the removal of Cd (II)

from its solution by the fungus, Rhizopus oligoorus, as chemically equilibrated

and saturated mechanism which reflected the predominantly site specific

mechanism on the cell surface. Drake et al., (1996) studied heavy metal

removal by using Datura innoxia. Veglio and Beolchini (1997) reviewed the

removal of heavy metal from waste solutions and the recovery of precious

metals by using various methodologies. Cervantes et al., 2001, presented a

review on the removal of Cr and its compounds by the use of living organisms,

such as bacteria, algae, fungi and plants. Various interactions and mechanisms

like biosorption, diminished accumulation, precipitation, reductions of Cr (VI)

15

(more toxic form) to Cr3+ (less toxic form) and chromate efflux were brought to

focus. Godlewska-Zylkiewicz (2006) reviewed the applications of bacteria,

yeast, algae and fungi for the pre-concentration of heavy metals from

environmental samples and the selective binding sites on the cell wall of these

groups functioning as sorbents of various heavy metals (Cd, Cr, Cu, Pb, Zn,

Fe, Au, etc.). Jianlong (2002) studied the biomass of the fungus,

Sacharomycete cerevisiae for the uptake of Cu and compared its function as

biosorbent by chemical modification of the components of cell wall. Vasudevan

et al., (2003) studied Cd biosorption in yeast. Davis et al., (2003a) studied the

physicochemical aspects of biosorption mechanism of toxic heavy metal, such

as Cd (II) by the use of algal biomass of Sargassum spp. Ahalya et al., (2005)

examined the husk of Bengal gram for adsorption of Cr (VI) and found that

carboxyl and hydroxyl groups were responsible for binding to metal ions. Haq

and Shakoori (1998) suggested that various organisms‘ ability to remove Cr

and bacteria reducing Cr (VI) (highly toxic hexavalent form) to Cr3+ (less toxic

trivalent form) can be exploited for the removal of toxic metals from wastewater.

Seiki and Suzuki (2006) advocated the phenomenon of biosorption and argued

that the strong metal-binding ability of biomass has attracted much attention in

the fields of wastewater treatment and environmental remediation

contaminated by toxic heavy metals. Iqbal et al., (2007) developed an

immobilized hybrid biosorbent for the sorption of Ni2+ from the aqueous

solution. McMahon et al., (2006) viewed microbial ecology as the potential to

transition from a purely descriptive to a predictive framework as regards their

capacity of bioremediation. Sari and Tuzen (2008) studied biosorption of Pb2+

and Cd (II) from aqueous solution using green alga, Ulva lactuva, and applied

various models, such as Langmuir, Freundlich and Dubinin-Radushkewich for

determining the equilibrium of biosorption. Volesky (2007) reviewed the

property of biomolecules on the surface of algal biomass responsible for

biosorption of heavy metals. Aydin et al., (2008) reported use of low cost

adsorbents derived from shells of wheat, rice and lentils for the removal of Cu2+

ions. Recently, kinetic and thermodynamic studies for the biosorption of lead

ion has been reported by Lawal et al., 2010. ‖

16

2.5. BIOSORPTION THROUGH THE USE OF ALGAE ―In addition to the use of synthetic adsorbents, such as fungi, bacteria,

yeast, etc., algal biomass has been tested for its potential to remove heavy

metal pollutants from the aquatic ecosystem. Chong and Volesky (1996)

studied metal uptake performance of a biosorbent prepared from a seaweed,

Ascophyllum nodosum, biomass using aqueous solutions containing Cu, Cd

and Zn ions in binary and tennary mixtures. Chong and Volesky (1996)

undertook biosorption studies using a sea weed, A. nodosum,biomass and

discussed equilibria of metal contents in tennary effluents. Leusch et al. (1995)

found that the ions of Pb and Cd could be bound very efficiently from very dilute

solutions by using the dried biomass of some ubiquitous brown marine algae,

such as Ascophyllum and Sargassum, which showed results in accumulation

and adsorption of metal ions in their dried biomass form. Figueria et al., (1996)

reported that nonliving biomass of Saragassum, a brown marine alga was

capable of binding more than 10% of its dry weight in toxic Cd ions. Gardea-

Torresdey et al., (1998) studied the ability of cyanobacteria to remove metal

ions from solution and demonstrated the presence of metallothionein gene in

various strains. Cook (1998) examined the surface layer of green algae, Nitella

gracilis, by using the technique of scanning electron microscopy (SEM) and

transmission electron microscopy (TEM). Kramer and Meisch (1999) introduced

carboxyl groups to the cell walls of fungal mycelia and studied the metal

binding abilities of the products for Cd (II), Co (II), Ni (II) and Zn (II) using to the

Langmuir model. Matheickal et al., (1999) studied Cd (II) adsorption by the

marine alga, Durvillea potatorum and reported that under pH 5.0, adsorption

was optimum. ‖

―Macfie and Welbourn (2000) focused on the polysaccharides and

glycoproteins of cell wall of algae and proposed the binding capacity of

negatively charged functional groups, such as carboxylic (COOH), sulfahydryl

(SH), etc. to heavy metal ions. They studied Cd, Co, Cu and Ni uptake by using

green algae, Chlamydomonas reinhand, and suggested that metal tolerance in

the alga must involve a complex of mechanisms involving both of external and

internal detoxification of metals. ‖

―Lee et al., (2000) investigated anion exchange type mechanism in the

uptake of chromium ions by using red, brown and green marine algae (48

17

species). Godlewska-Zylkiewicz (2006) reviewed the applications of bacteria,

yeast, algae, and fungi for the preconcentration of heavy metals from

environmental samples and the selective binding sites on the cell wall of these

groups functioning as sorbents of various heavy metals (Cd, Cr, Cu, Pb, Zn,

Fe, Au, etc.). Rehman and Shakoori (2001) carried out studies on metal

resistance alga, Chlorella spp., and removal of heavy metal like Cr (VI) from the

environment by the use of algal strains involving the mechanisms of

biosorption, adsorption, and bioaccumulatiuon. Volesky (2001) declared the

biomass of dead algae as a potential technology for toxic metal removal and

discussed its reuse and cost-effectiveness. Cossich et al., (2002)

recommended that treating the industrial effluent with the biomass of the

seaweed, Sargassum sp., is advantageous in high metal (Cr) binding capacity

as well as low cost.

Davis et al., (2003a) reviewed the removal of toxic heavy metals, such

as Cd (II), Cu (II), Zn (II), Pb (II), Cr (III) and Hg (II) by the use of brown alga

and related the polysaccharides of its cell wall containig their key functional

groups responsible for adsorption. Davis et al., (2003a) studied the metal

selectivity of saragassum spp. and their alginates in relation to their glucuronic

acid content and conformation. Carmona et al., (2005) studied the removal of

toxic Cr3+ and Cr (VI) from wastewater in a factorial experimental design using

Saragassum as biosorbent. Chojnacka et al., (2005) studied biosorption,

desorption and binding sites of heavy metals including Cr (III), Cd (II) and Cu

(II) on the cell wall of blue-green algae, Spirulina sp. Rao et al., (2005) worked

on biosorption of Ni and Cu from aquous solution by using dead biomass of

green algae, Sphaeroplea. Mehta and Gaur (2005) reviewed the biosorption of

toxic heavy metals by the use of algae by focusing on the binding sites

(carboxyl group) on the cell surface of the algae. Naja et al., (2005) studied the

removal and recovery of heavy metals from solutions using physico-chemical

mechanisms including biosorption, precipitation and microbial reductive

processes. The biomass of a brown seaweed, Sargassum fluitans was

documented for metal biosorption of Cu2+

, Zn2+

and Cd (II). Carboxylic,

sulfonate and phosphonate moieties of the biomass were confirmed

quantitatively to be involved in the uptake process of heavy metals. Naja and

Volesky (2009) studied ion exchange mechanism of biosorption of heavy

18

metals, such as Cu2+, Zn2+ and Cd (II) by the dead biomass of S. fluitans.

Romera et al.,(2006) reviewed biorption of various toxic metals such as Cd

(II),Cu2+, Ni2+, Pb2+ and Zn2+ using 37 different algae (20 brown algae, 9 red

algae and 8 green algae) and declared that brown algae stand out as very

good biosorbents and statistically analyzed the biochemical affinities. Vilar et

al., (2005) studied algal waste from agar extraction for the removal of Cd (II)

from aquous solution with special emphasis on equilibrium and kinetics.

Volesky (2007) reviewed the property of biomolecules on the surface of algal

biomass responsible for biosorption of heavy metals. Pahlvanzadeh et al.,

(2010) reported the equilibrium, dynamic and thermodynamic studies of

biosorption of Ni+2 by using brown algae. ‖

2.6. GREEN ALGAE AND BIOSORPTION

―Green algae (Chlorophyceae) have been largely ignored in the early

years of research on heavy metal biosorption studies; however, recently, a

number of studies have been reported in this context. Donmez et al., (1999)

tested three algal species, Chlorella vulgaris, Scenedesmus obliquous and

Sanechocystis sp. for the biosorption of Cu2+, Ni2+ and Cr (VI) as a function of

pH and found adsorption models of Langmuir and Freundlich suitable for

describing the results. Ozer et al., (1999) studied the adsorption of Fe, Pb and

Cd ions onto green algae, Schizomeris leibleini and suggested optimium pH of

2.5, 4.5, and 5.0 for all these metals respectively, at 30°C. The experimental

data were in agreement with both Langmuir and Freundlich isotherms.

Adsorption of Zn2+ by Cladophora crispata was examined by Ozer et al., (2000)

recommending the removal of heavy metal from wastes by using dried algae.

Lee et al., (2000) investigated anion exchange type mechanism in the uptake of

chromium ions by using various types of algae including the ones belonging to

green algae. Gupta et al., (2001) declared Spirogyra species suitable for the

removal and recovery of Cr (VI) from wastewater. Tang et al., (2002), predicted

equilibrium model for Cd adsorption by green algae in a batch reactor. The

models for Cd adsorption were found to be applicable over a pH range of 4.5–

10.5, even though it was found that pH affects the Cd adsorption potential

significantly. Terry and Stone (2002) studied removal of Cd and Cu from water

via biosorption using the green algae.

19

Akhtar et al., (2003) developed biosorbent for the removal of heavy metal by

immobilizing a unicellular green microalga Chlorella sorokiana within luffa

songe discs. Cruz et al., (2004) investigated Cd biosorption by using

Saragassum. ‖

―Kola et al., (2004) investigated green alga, Chlamydomonas reinharditi

for its ability to sequester and detoxify heavy metals, such as Cd by focusing on

cadmium binding sites of cel wall. Rao et al.,(2005) studied the biosorption of

Ni2+ and Cu2+ using the dead alga, sphaeroplea. Shuja and Azzizullah (2006)

reported the effect of biomass concentration on the extent of biosorption of

chromium by using Spirogyra Sp.‖ ‖

―Andrade et al., (2005) focused on binding sites on the cell wall of green

alga Chaetophora elegans, for adsorption of heavy metals including Cd (II),

Pb2+, Ni2+ and Zn2+. Arica and Bayramoglu (2005) investigated biosorption of Cr

(VI) by Chlamydomonas reinhardtii. ‖

―Frailes et al., (2005) studied the sorption capacity of microalga,

Chlorella vulgaris for varying concentrations of Cu, Zn, Cd and Ni. Tu zu n et al.,

(2005) screened microalga, Chlamydomonas reinhardtii for biosorption of Hg2+,

Cd (II) and Pb2+ ions and found the involvement of amino, carboxyl, hydroxyl

and carbonyl groups on its cell wall. Doshi, (2006) elucidated investigations on

ability of Chlorella sp. to adsorb metal ions like Cu2+ and Ni2+ by using infrared

and SEM. Jagiello et al., (2006) studied Spirulina for the adsorption of Cr3+. ‖

Romera et al.,(2006) analyzed biorptive potential of various toxic metals such

as Cd (II),Cu2+, Ni2+, Pb2+ and Zn2+ using 37 different alga including 8 different

species of green algae, evaluated their biosorptive potential and analyzed their

biochemical affinities to heavy metal ions. Shuja and Azizullah (2006)

investigated the effect of pH on the capacity of biosorption of nickel by using

green algae, Oedogonium sp. Doshi, (2006) carried out spectroscopic and

kinetic studies for the sorption of Cd (II) in live and dead alga Spirulina

Sp.‖Gupta and Rastogi (2008) carried out studies on kinetic and equilibrium

aspectrs of biosorption of lead by using Spirogyra species.

2.7. ADSORPTION ISORTHERMS

―Saha et al., (2002) evaluated various adsorption isotherms such as

Langmuir isotherms, Freundlich isotherms, etc. and suggested langmuir

20

isotherms to be the best in describing the adsorption data of phosphate by

hydroxyl inter-layered expansible clays. Ozer et al., (1999) found that

experimental biosorption equilibrium data for Fe, Pb and Cd ions were in

agreement with the dataas calculated by Freundlich and Langmuir models.

Various adsorption isotherms were applied to biosorption of heavy metal ions

by algal biomass and application Langmuir isotherms and Freundlich isotherms

were found to yield best results. Wang and Chen (2006) predicted heavy metal

biosorption by biosorbents on the basis of the kinetic studies. Okuo et al.,

(2006) reported biosorption of selected heavy metal ions and fitted the data to

various adsorption models, such as Langmuir, Freundlich, etc. Aravindhan et

al., (2004) elucidated the mechanism of the process of biosorption by applying

various models, such as first order and second order kinetic models, etc. Igwe

and Abia (2007) compared various isotherms regarding adsorption of heavy

metal and found that Dubinin-Radushkevich gave the best fit with R2 values

ranging from 0.9539 to 0.9973 and an average ―R2― value of 0.9819, followed

by Freundlich isotherm (average ―R2― values = 0.9783) and then the Langmuir

isotherm (Ave. 0.7637). Currently, a trend has emerged to evaluate biosorption

performance by determining equilibrium, kinetic and thermodynamic

parameters. A number of biosorption studies have been reported by using

various biosrbents in this context (Gupta and rastogi, 2008; Ghadbane et al.,

2008; Pahalvanzade et al., 2010; Lawal at al., 2010). ‖

2.8. CELL WALL COMPOSITION AND BINDING SITES

―Biochemical agents (binding sites), such as cellulose and

peptidoglycans on the cell wall of various biosorbents, such as algae

(microalgae, green algae, red algae, brown algae, etc.), fungi and bacteria

have been found to play a key role in the adsorption of heavy metal cations.

Wurdack (1931) determined the composition of the cell walls of some green

algae, including Cladophora, Spirogyra, Oedogonium, etc. Dawes (1965)

studied Spirogyra sp. and focused on the outer amorphous layer of cell wall

covering cellulose. Thompson and Preston (1967) suggested the presence of

proteins bound to the cell wall of algae. Grant (1979) studied the specific

binding sites of divalent cations to polysaccharides, leading to the

understanding of the possibility of cell wall components‘ binding to metal ions.

Evans (1974) studied the synthesis and composition of extra cellular mucilage

21

in the unicellular red alga Rhodella. McDnald and Heaps (1976) examined the

ultrastructure and differentiation in Cladophora glomerata. Crist et al., (1981)

examined the metallic ions adsorbed by the algal cell walls and the nature of

bonding involved. Becker et al., (1988) investigated the cell wall of Tetraselmis

striata (Chlorophyta) and focused on macromolecular composition and

structural elements of the complex polysaccharides. Cook (1998) examined the

surface layer of the green alga, Nitella gracilis by using the technique of

scanning electron microscopy (SEM) and transmission electron microscopy

(TEM). ‖

―Kramer and Meisch (1999) introduced carboxyl groups to the cell walls

of fungal mycelia and studied the metal binding abilities of the products for Cd

(II), Co2+, Ni2+ and Zn2+ according to the Langmuir model. Fleurence (1999)

elucidated studies in the enzymatic degradation of cell wall polysaccharides of

seaweeds belonging to the groups of red algae and green algae. Macfei and

Welbourn (2000) studied ultrastructure and chemistry of the cell wall of red

algae, Chlamydomonas reinhardtii (Chlorophyceae). Davis et al., (2003 a)

focused the linear polysaccharides on the cell wall of Saragassum fluitants and

Saragassum siliquosum explained the metal biosorption of metal ions. Deniaud

et al., (2009) developed a protocol for physical, chemical and enzymatic

treatment of cell wall of red alga, Palmaria palmaria to prepare the fractions

enriched in proteins. Raize et al., (2004) investigated the involvement of

chemical groups on the cell wall in mechanism of biosorption of metallic cations

(Cd, Ni, Pb). Arief et al., (2008) reviewed a number of investigations regarding

binding sites of biosorption on the surface of algae. ‖

2.9. MECHANISM OF BIOSORPTION

―Drake et al., (1996) studied a shrubby plant, Datura innoxia at cellular

level and characterized metal ion binding sites on it by using lanthanoid ion

probe analysis. Cook (1998) indicated the binding site of heavy metlas on the

surface layer of green alga, Nitella gracilis by using the technique of scanning

electron microscopy (SEM) and transmission electron microscopy (TEM).

Kratochvil and Volesky (1998) suggested that various functional groups, such

as carboxyl and sulfate adsorb to heavy metal by ion exchange mechanism.

Tereshina et al., (1999) focused on the mechanism of adsorption of heavy

22

metal ions by cell wall polysaccharides of the micro-alga, Aspergillus niger.

Whiffen et al., (2007) applied field emission scanning electron microscopy

(FESEM) technique to study the ulltrastructure of algae, Mougeotia. Cavalcante

et al., (2004), investigated chromium adsorption by using the marine algae

Saragassum sp. in a column of fixed bed. Davis et al., (2003b) reported the role

of alginate on the cell wall of Saragassum sp. for the removal of heavy metal

cations. Slaveykova et al., (2003) presented a review study on the biotic ligand

model to appreciate mechanistic understanding of the interactions of metal ions

with biological surfaces. Naja et al., (2005) studied the acidic functional groups

on the cell wall of bacteria and fungi for the binding of Pb. Alscher et al., (2002)

studied the characteristics of certain carbohydrates with respect to metal

binding and the modification in its structure which could enhance its affinity with

metal ions. ‖

―Rosen (2006) declared that the modern approaches like liquid

chromatography and tandem mass spectrometry, as increasingly becoming a

key technique for environmental analysis. Sankarramakrishnan et al., (2007)

used chitosan derivative for the removal of cadmium ions from electroplating

waste effluents under laboratory conditions. Yahya et al., 2009 reported the

mechanism of biosorption of Cu2+ onto the immobilized cells of Pycnoporus

sanguineus. Lawal et al. (2010) elucidated the mechanistic aspects of the

process of biosorption. ‖

―Considering the growing understanding of biosorptive potential of green

algae and the possible binding sites on its cell wall, utilization/exploitation of

filamentous green algae is all the more significant in order to appreciate its

performance in biosorption of toxic heavy metals and to introduce improved,

efficient and cost-effective biosorbents in the exponentially increasing issue of

wastewater treatment. ‖

23

CHAPTER 3 MATERIALS AND METHODS

3.1. MATERIALS

3.1.1 BIOSORBENTS

―Selected species of filamentous green algae, used for the biosorbent

study, were tested for their biosorptive capacity for heavy metals, such as

cadmium (Cd) and Chromium (Cr) ions (sorbate). The biosorbents

(biomass of dead algae) employed in this study were acquired from the

following sources:‖

Spirogyra communis (Hull et al., 1985) was obtained from the

University of Texas, Culture Centre, UTEX culture no., LB

2466, USA. Pure algal cultures were raised and maintained at

the Aquatic Entomology and Ecology Laboratory, Mid Florida

Research and Education Center (MREC), University of Florida,

USA.

Cladophora delmatica (Schneider and Searles, 1991) was

collected from a natural wetland located at 28º 24' 40" N and

81º 33' 55" W in central Florida, USA, and was taken to the

laboratory at MREC, University of Florida for subsequent

studies.

Spirogyra spp., mixture belonging to three different species, viz:

S. juergensii, S. elongate, and S. piepengensis were collected

in polythene bag from Botanical Gardens of GC University, The

Mall, Lahore, Pakistan.

Culture conditions of Spirogyra communis

―Spirogyra communis was raised at the Aquatic Entomology and Ecology

Laboratory, MREC, University of Florida. The Cultures were grown in Bold 1NV

Medium (Table:3.1 ), in 1 Liter Erlenmeyer flask at 20oC, under cool

fluorescent lamps at maximum intensity of 3200 lux and periodicity of 12/12

24

hour Light/Dark. The growth conditions were optimized, and medium was

periodically refreshed by changing the portion of the medium. ‖

Pretreatment of Biomass

―The biomass was thoroughly washed with distilled water to remove all

the extraneous material and placed on a filter paper to reduce the water

content prior to treating the biomass with 0.02 M HNO3. It was then dried

overnight at 50o C until a constant weight was achieved and the final weight of

the biosorbent was recorded. The biosorbents were then crushed and passed

through a 300 nm sieve to obtain uniform particle size of each biosorbent used

for further studies. The same procedure was adopted for biomass of all

experimental algae S. communis, C. delmatica, and Spirogyra spp.‖

3.1.2. PREPARATION OF HEAVY METAL [Cd (II) AND Cr (VI)] IONS

SOLUTIONS

3.1.2.1. Cadmium Solution

―For Cd (II), a stock solution of Cadmium nitrate (CdNO3) was prepared

by dissolving 27.0 grams of CdNO3 (Analytical grade, Fisher Scientific, USA) in

100 mL of double distilled water to make a concentration of 1000 mg/L, and

serial dilutions from of this stock solution were prepared to obtain10, 20, 30, 40,

50, and 100 mg/L concentration of Cd (II) ion solution.‖

3.1.2.2 Chromium Solution

―For Cr (VI), a stock solution of Potassium Dichromate (K2Cr2O7) was

prepared by dissolving 83.0 grams of K2Cr2O7 (Analytical grade, Fisher

Scientific, USA) in 100 mL of distilled deionized water to make a concentration

of 1000 mg/L, and from this stock solution, serial dilutions were made to obtain

100, 200, 300, 400, and 500 mg/L concentrations of Cr.‖

25

3.2. METHODS AND PROCEDURES

3.2.1. BISORPTION STUDIES

3.2.1.1. Glassware and Apparatus

―All biosorption experiments related to biosorptive potential of algal

biomass were conducted by using 125 mL Erlenmeyer flasks. Prior to use, the

flasks were baked at 70°C for 4 hours, followed by one wash with concentrated

HNO3 and then one wash with distilled water. ‖

3.2.1.2. Biosorption Evaluation

―One hundred mL of Cd (II) and Cr (IV) ion solutions in 125 mL

Erlenmeyer flasks were taken; Cd (II) ion solutions at concentrations of 10, 20,

30, 40, 50 and 100 mg/L were taken while that of Cr (VI) ion solutions were

taken100, 200, 300, 400, and 500 mg/L concentrations. The values of pH of all

the solutions were monitored by pH meter throughout the experiment and

adjusted according to the experiment by using 0.2 N HNO3 and 0.1 N NaOH.

―Spirogyra communis, C. delmatica, and Spirogyra spp., each in the

amount of 0.1, 0.2, and 0.3 g (1g/L, 2g/L, and 3g/L respectively) of dried algal

biomass were introduced in the flasks of the above mentioned Cd (II) and Cr

(VI) ion concentrations separately. All the three biosorbents under investigation

were also introduced to flasks filled with pure distilled water with no metal ion

(control). The flasks were maintained at 25ºC under constant agitation on a

rotator shaker (200 rpm) for a period of 3 hours.‖

―The values of pH for the biosorption capacity of algal biomass I, II and

III for Cd (II) and Cr (VI) were optimized by a series of initial experiments.

Finally, studies for Cd (II) ions biosorption were carried out at pH value of 5.0

and those for Cr (VI) ions biosorption at 4.0 for all the subsequent studies,

unless otherwise mentioned. The optimum sorbate (metal ions) and biomass

(algal biomass) were also optimized and finally, subsequent studies relating to

26

Cd (II) ion solutions were performed at 100 mg/L of sorbate concentrations and

those of Cr (VI) ion solutions at 400 mg/L . The biosorbents (biomass of dead

algae) under investigation were used in the concentration of 1 g/L unless

otherwise mentioned. All the studies were carried out at 25°C unless otherwise

mentioned.‖

All the experiments were conducted in triplicate.

3.2.1.3. Transport of Samples for Analysis

―Samples of 5 mL from each flask were collected for subsequent

determination of residual concentrations of metal ions at regular intervals of 0,

30, 60, 120, 150 and 180 min. All the samples were then passed through

Osmonics/MSI* Cameo* Glass/Nylon Syringe Filters 0.25 nm (Fisher Scientific,

USA). The adsorption capacity of the filter for Cd (II) and Cr (VI) ions had

already been tested which was less than 5% for both of the ions. The filtrate

was then preserved for further analysis. ‖

3.2.1.4. Analysis by Atomic Absorption Spectrometer

―All the samples were tested for metal ion concentration by using Atomic

Absorption Spectrometer at University of Florida‘s Analytical Research

Laboratory, Gainesville, USA, for the analysis. A few samples were also tested

by using Atomic Absorption Spectrometer facility, at theDepartment of

Chemistry, of University of the Punjab, Lahore, Pakistan. ‖

3.2.2. EVALUATION OF BIOSORPTIVE BINDING SITES

3.2.2.1. Fourier Infra-Red Transform Spectroscopy (FTIR)

―For FTIR spectroscopy, tablets of algal biomass were prepared in a

Graseby-Specac Press, using algal mass mixed with Potassium Bromide (KBr,

1:100 p/p). The following samples were subjected to FTIR. ‖

Algal biomass cleaned by HNO3 treatment and deionized water (control)

Cadmium loaded (Cd (II) - loaded) algal biomass

Chromium loaded (Cr (VI) - loaded algal biomass

―FTIR Spectra of biosorbents under investigation [Cd (II) ions treated, Cr

(VI) ions treated, and non-treated (Control)] were obtained at a resolution of 1

cm-1. A window between 500 and 4500 cm-1 for C. delmatica and Spirogyra

27

spp., studies, and 1000 and 4000 cm-1 for S. communis. The above mentioned

range of spectra contains information of probable characteristic

polysaccharides which will be responsible for the modification of cell wall

structure after being treated with Cd (II) and Cr (VI) ions. All spectra were

normalized and baseline corrected with Perkin-Elmer IR Data management

software. Data were then exported to Microsoft Excel 2003 and all spectra were

area normalized.‖

These analytical studies were performed at the Department of

Chemistry, University of the Punjab, Lahore, Pakistan, and Pakistan Council of

Scientific and Industrial Research (PCSIR), Lahore, Pakistan.

3.2.2.2. Scanning Electron Microscopy (SEM):

―The microscopic studies of the surface of biosorbents under

investigation [Cd (II) ions treated, Cr (VI) treated, and Non-treated (control),

exposed to pure distilled water] were carried out by Scanning Electron

Microscopy (SEM). Preparation of the samples for SEM studies was carried out

by following the protocol as suggested by Cook, (1998). Specimens were fixed

in a solution of one part Karnovsky's fixative (4% paraformaldehyde and 0.5%

glutaraldehyde), one part 2% osmium, and one part culture medium [Bold N1

medium, UTEX) for 10 min, followed by six 4-min rinses in distilled water.

Specimens were left in rinse water overnight, dehydrated in a graded ethanol

series, dried, using a Tousimis Samdri-780-A critical-point drier, mounted on

aluminum stubs with double-sided tape, sputter-coated using a Bio-Rad

E5000M gold coated, and viewed using a Hitachi S570 scanning electron

microscope at either 10 or 20 kV. This procedure was performed by using the

appropriate facility at University of Burdwan, West Bengal, India. ‖

‖

3.2.3 MATHEMATICAL MODELING AND INTERPRETATION OF DATA

―The data collected as a result of biosorption studies were tested by

conventionally used adsorption isotherms, such as Langmuir and Freundlich

isotherms for adsorption. Moreover, kinetic modeling and thermodynamic

studies were also carried out. ‖

28

3.2.3.1 Equilibrium Isotherm Modeling

―Langmuir isotherms were used to correlate the equilibrium data.

Langmuir model assumes a monolayer sorption of sorbate from the aqueous

solution (Mashilah et al., 2008). The Langmuir equation is given below

(Langmuir, 1918):

Here,

qe = Equilibrium constant of sorbate ion on surface of the biosorbent (mg/g)

Ce = Equilibrium concentration of metal ion in solution

b = Langmuir‘s constant (L/mg)

― The values of ―qe‖, ―Ce‖, and ―b‖ were calculated from intercept and slope

of linear plot of ―1/qe‖ versus ―1/ Ce‖. The distribution coefficient (k) for metal

ions between the sorbent and the aqueous solution at equilibrium stage was

determined from the following expression:‖

―The Freundlich model (Freundlich, 1906) was also employed to

estimate the adsorption intensity of the adsorbent towards the sorbate. This

theorem considers multi-layers adsorption on the sorbent surface. This model

can be explainedby the equation below:‖

=

Here,

Kf = Freundlich emperical constant relative to sorption capacity

1/n = emperical constant relative sorption Intensity

The values of the ―Kf‖ and ―n‖ were derived from the intercept and slope

respectively of a linear plot of ―lnqe‖ versus ―lnCe‖.

Temkin isotherm was also employed as given by the following equation (Wang

and Qin, 2003):

―KT‖ = equilibrium binding constant correlated to the maximum binding energy

and

29

―B‖ = constant related to the heat of adsorption.

A linear plot of ―qe‖ verses ―ln Ce‖ enables the determination of the isotherm

constants, ―B‖ and ―KT‖ from the slope and the intercept respectively.

3.2.3.2 Kinetic Studies

―Pseudo-second order kinetic model as developed by Ho and McKay

(1999) was employed to evaluate the kinetic parameters for the biosorption

studies of Cd (II) and Cr (VI) ions by biosorption from dead biomass of S.

communis, was investigated. The concentration of biomass was kept at 1 g/L

while initial Cd (II) ion concentration was varied from 10 to 40 mg/L and that of

Cr (VI) varied from 100 to 400 mg/L.‖

―To explain the correlation between the equilibrium concentration of

metal ions in the solid phase (sorbent) and the aqueous solution, pseudo-

second order model, as developed by Ho and McKay (1999), was used.

Pseudo-second-order model considers that the rate of occupation of

biosorption sites is proportional to the square of the number of unoccupied

sites. ‖

Where,

t (min) = time

qt (mg g−1) = uptake capacity at time ‗t‘

K2 (gmg−1 min−1) = equilibrium rate constant of pseudo-second-order

adsorption. After being integrated and rearranged, following expression could

be achieved:

The values of the constants were calculated by plotting ―t/qt‖ versus ―1/qe‖.

3.2.3.3 Thermodynamic parameters

―Thermodynamic parameters were also determined form the

experimental data. Van‘t Hoff equation (Pahlvanzadeh et al., 2010) was used to

obtain the values of entropy (∆H) and enthalpy (∆S). ‖

30

―Where ―R‖ is gas constant; its value is equal to 8.314, and ―b‖ is the

Langmuir‘s constant. The linear plot of the ―lnb‖ versus ―1/T‖ (Figure 4.9) was

drawn and by determining the intercept and slope, values of entropy (∆H) and

enthalpy (∆S), were measured.‖

―Arhenius equation was used to obtain the values of Arhenius constant

(A₀) and the activation energy (Ea) by using the following equation.

The values of Gibbs free energy (∆G) for biosorption were determined by using

the equation given below (Pahlvanzadeh et al., 2010):

31

CHAPTER 4

RESULTS

―4.1. BIOSORPTION CAPACITY OF CADMIUM BY

Siprogyra communis

―The efficiency of the biosorption is strongly affected by the physico-

chemical characteristics of the solutions, such as pH, temperature, initial

concentration of the pollutants, etc. (Volesky, 2007). The biosorption capacity

of S. communis was studied under variable conditions. ‖

4.1.1.1. Effect of pH

―The biosorption of Cd (II) was investigated at pH values of 1.0, 2.0, 3.0,

4.0, 5.0, 6.0 and 7.0 (biosorbent concentration at 1g/L; initial sorbate

concentration at 100 mg/L). Biosorption of Cd (II) ions was very low at pH value

of 1.0, increasing gradually with the increase with the increase in pH. The

maximum biosorption of 43 mg/g was achieved at pH 5.0, beyond this pH value

biosorption capacity declined. ‖

4.1.1.2. Effect of Contact Time

―The biosorbent exposure in terms of contact time for the sorbate was of

importance as evident from Fig. 4.2 showing the biosorption efficiency of Cd (II)

by S. communis as a function of contact time (biosorbent: 1g/L; initial sorbate

concentration: 100 mg/L; pH: 5.0). The efficiency of biosorption increased

significantly with the increase in contact time during the first 30 min., followed

by slower uptake up to 120 min. of contact time and thereafter, equilibrium was

probably achieved because no significant uptake of Cd (II) ions after 120 min.,

of exposure.‖

4.1.1.3. Effect of Temperature

―Temperature is very important kinetic parameter in the biosorption

process. It affects the mobility of sorbate ions as well as biosorption capacity of

the biosorbent. In this study, effect of temperature on the biosorptive capacity

has been studied at 10oC, 20oC, 30oC, and 40oC (biosorbent: 1g/L; initial

32

sorbate concentration: 100 mg/L; pH: 5.0) revealed that the extent of

adsorption was maximum at 20oC (43 mg/g) followed by that at 30oC (Fig. 4.3),

with lower adsorption noticed at the temperatures 10oC and 40oC. Temperature

also affects the time required to achieve maximum biosorption (qmax) as was

observed in this study that the time required to achieve ―qmax‖ was reduced at

elevated temperatures.‖

4.1.1.4. Effect of Biosorbent Quantity

―The effect of biomass dosage on the biosorption of Cd (II) ions at

different biomass concentrations of 1 g/L, 2 g/L, and 3 g/L (initial sorbate

concentration: 100 mg/L; pH: 5.0) showed that the biosorption efficiency is

highly dependent on the increase in biomass dosage of the solution (Fig. 4.4).

The biosorption of Cd (II) by S. communis was found to be directly proportional

to the biomass concentration. However, at exposure of low sorbate

concentration of 1 mg/L of Cd (II) ion solution, biosorption of Cd (II) ions per

gram biosorbent (qe) was maximum, achieving a value of 43 mg/g, whereas at

3 g/L dose, it was only 18 mg/g (Fig. 4.8). ‖

4.1.1.5. Biosorption Equilibrium Isotherm

―Various sorption models were employed for fitting the data to examine

the relationship between sorption and aqueous concentrations of metal ions

[Cd (II)]. This study was performed at pH value of 5.0, temperature 25°C

biosorbent concentration of 1 g/L, temperature of 25°C, and varying sorbate

concentrations of 10, 20, 30, 40 mg/L.‖

4.1.1.5.1. Langmuir Isotherm

―Langmuir isotherm was employed to correlate the equilibrium

concentrations of Cd (II) ions onto the biosorbent, equilibrium metal ion uptake

capacity (qe) and equilibrium metal ion concentrations in the solution (Ce).

Different ―qe‖ and ―Ce‖ were determined for different initial concentrations.

These values were then used to draw Langmuir isotherm plot according to

following equation.

33

The linear form of the above equation may be written as:

― The plot of inverse of equilibrium concentrations of Cd (II) ions uptake by

biosorbent, ―qe‖ (mg/g), versus inverse of equilibrium concentrations of Cd (II) in

the aqueous solution, ―Ce‖ (mg/L), is shown in Fig. 4.5. A typical equilibrium

biosorption isotherm is obvious from figure, suggesting that biosorption of Cd(II)

ions involves a chemical equilibrated and saturable mechanism which reflects

site-specific biosorption on the surface of the sorbent. Langmuir‘s maximum

uptake capacity (qmax) of S. communis for Cd (II) ions was found to be 1.44

mg/g, the value of saturation constant (b) was calculated as 2.5 Lmg-1, and the

regression coefficient (R2) value was calculated as 0.9605. The value of

distribution coefficient (k) for Cd (II) ions between the adsorbent and the

aqueous solution (qe/Ce) was 0.75 (Table 4.1). ‖

4.1.1.5.2. Freundlich Isotherm

―Freundlich model was employed to estimate the adsorption intensity of

the adsorbent towards the sorbate [Cd (II)] as given by the equation below‖

=

The linearized form of the above equation is given as:

―Where, ―KF‖ is Freundlich‘s constant relative to sorption capacity, and

―n‖ is the emperical constant relative sorption Intensity. Fig. 4.6 shows a linear

relationship of Log ―qe‖ versus Log‖ Ce‖. The value of Freundlich‘s saturation

constant (KF) was found to be 14.49, constant (n) was 0.16 (derived from

intercept and slope respectively), and regression coefficient value was (R2) was

0.9986 (Table 4.2). ‖

4.1.1.5.3. Temkin isotherm

Temkin model can be presented as

The linearized form of the above equation 4.5 can be given as:

34

―Where, ―KT‖ is Temkin‘s equilibrium binding constant corresponding to

maximum binding energy, and ―B‖ is Temkin‘s constant related to heat of

adsorption. Linear plot of ―qe‖ versus ―Ce

‖ was drawn (Fig. 4.7) and the values of

―KT‖ was calculated as 0.024, the constant, ―B‖ was 4.31 (derived from the

slope and the intercept values, respectively), and the regression coefficient

very highly significant (R2 = 0.9999) (Table 4.3). ‖

4.1.1.6. Kinetic Studies

―Pseudo-second order kinetics of the uptake of Cd (II) ions by dead

biomass of S. communis, was investigated (biomass concentration: 1 g/L,

pH=5.0, and initial sorbate concentrations: 10 mg/L). To explain the correlation

between the equilibrium concentration of metal ions in the solid phase (sorbent)

and the aqueous solution, following model, as developed by Ho and Mckay

(1999), was used. ‖

―Where, ―qt‖ is the uptake of heavy metal ion (mg/L) by the biosorbent at

any time

(t), ―qe‖ is heavy metal ion concentration (mg/g) in the biosorbent at equilibrium

stage, ―K2‖ is pseudo-second order rate constant (g.mg-1.min-1), and ―t‖ is time

(min) interval of sampling.

―From the linear plot of ―t/q‖ versus ―t‖ (Fig. 4.8), the value of ―qe‖ (Cal)

was found to be 43.61 (slope of the graph), ―K2‖ for the adsorption mounted to

0.10 g.mg-1.min-1 (derived from intercept of the plot), and the ―R2‖ value was

0.9999 (Table 4.4).‖

4.1.1.7 Thermodynamics studies

―Thermodynamic behavior of the adsorption of Cd (II) ions on the

surface of S. communis was studied by calculating various thermodynamics

35

constants. The following equation was used to obtain the values of entropy

(∆H) and enthalpy (∆S):‖

―Where, ―R‖ is gas constant, its value is equal to 8.314 and b is the

Langmuir‘s constant. The linear plot of the lnb Versus 1/T (Fig. 4.9) was drawn

and by determining the intercept and slope, values of entropy (∆H) was

calculated as 8.50 J/mol and enthalpy (∆S) as 41.8 J mol-1K-1 (Table 4.5). In

addition, Arhenius equation was used to draw a linear plot of ―ln K2‖ versus

―1/T‖ (Fig. 4.10) and to obtain the values of Arhenius constant (A₀) and the

activation energy (Ea) as 4.2 and 35.58 J mol-1g-1, respectively (Table 4.6) by

using the equation given below. ‖

―The values of free energy (∆G) for sorption of Cd (II) ions on S.

communis were calculated as 12.31, 20.03, 21.95, and 11.43 at temperatures,

283 K°, 293 K°, 303 K°, and 313 K°, respectively (Table 4.7), by using the

following equation: ‖

4.1.2. BIOSORPTION CAPACITY OF CHROMIUM BY S.

communis


―The biosorption of Cr (VI) was studied at varying pH values, such as

1.0, 2.0, 3.0, 4.0, 5.0, 6.0 and 7.0 pH units (initial sorbate concentration of 400

mg/L; sorbent concentration of 1g/L). The optimum adsorption of Cr (VI) was

recorded at a pH value of 4.0 (242 mg/g) (Fig.4.11). With the declining pH

(˂4.0), decrease in Cr (VI) biosorption was observed. Whereas, with the

increase in pH above 4.0, decrease of biosorption was observed.


―Contact time of the exposure of biosorbent to the sorbate was observed

as an important factor in biosorption application. Fig. 4.12 shows that the

biosorption efficiency of Cr (VI) by S. communis as a function of contact time

36

(initial sorbate concentration: 400 mg/L; sorbent concentration: 1g/L). Rapid

uptake of Cr (VI) was observed in the first 30 min. (200 mg/g), followed by

relatively slower uptake up to 120 min. (260 mg/g); and thereafter, no

significant uptake of Cr (VI) ions was observed.


―Effect of temperature on the biosorptive capacity was studied at varying

temperatures of 10oC, 20oC, 30oC, and 40oC (initial sorbate concentration: 400

mg/L; sorbent concentration: 1g/L). The extent of adsorption was found to be

maximum at 20oC (264 mg/g), followed by 30oC (Fig. 4.13). However, lower

adsorption was found at the temperatures of 10oC and 40oC. ‖


―The effect of biomass dosage on the biosorption of Cr (VI) ions was

studied by using different biomass concentrations of 1g/L, 2g/L, and 3g/L (Initial

sorbate concentration: 400 mg/L; pH: 4.0). Results showed that the biosorption

efficiency is highly dependent on the increase in biomass dosage of the

solution (Fig. 4.14). The biosorption capacity of Cr (VI) by S. communis was

found to be directly proportional to the biomass concentration. However, at

1mg/L biosorbent concentration, maximum Cr(VI) ion uptake per unit volume

was observed (264 mg/g) as evident from Fig. 4.14. ‖


―Various sorption models are widely employed for fitting the data to

examine the relationship between sorption and aqueous concentrations of

metal ions [Cr (VI)]. This study was performed at pH: 4.0; biosorbent

concentration was 1 g/L, and different concentrations of Cr (VI) ion solutions

were 100, 200, 300, 400 mg/L.‖


―Langmuir isotherms were used to correlate the equilibrium data by

using the equation described previously. The values of ―1/qe‖ and ―1/Ce‖ were

derived from the biosorption experiments performed at different initial Cr (VI)

37

ions concentration, such as 100 mg/L, 200 mg/L, 300 mg/L, and 400 mg/L in

solution and linear plot of ―1/qe‖, versus ―1/Ce‖ of Cr (VI) was drawn as shown in

Fig. 4.15. This figure shows a typical equilibrium biosorption isotherm,

suggesting that biosorption of Cr (VI) ions involves a chemical equilibrated and

saturable mechanism which reflects site-specific biosorption on the surface of

the sorbent. Langmuir‘s ―qmax‖ for biosorption of as 411 mg/g; the value of

saturation constant (b) was calculated as 0.01 L/mg, and the regression

coefficient was highly significant (R2 = 0.9998) (Table 4.1). The value of

distribution coefficient (k) for Cr (VI) ions between the adsorbent and the

aqueous solution (qe/ce) was found to be 1.08 (Table 4.1).‖


―Freundlich isotherm model as previously described was employed to

estimate the adsorption intensity of the adsorbent towards the sorbate [Cr (VI)];

a linear relationship of ―Log qe‖ versus ―Log Ce‖ (Fig. 4.16). The values of

Freundlich‘s saturation constant (KF) was found to be 27.38, n was 0.85

(derived from intercept and slope respectively), and value of ―R2‖ was 0.9988

(Table 4.2). ‖


―Temkin isotherm as previosly described was used and the values of ―KT‖

(equilibrium binding constant corresponding to maximum binding energy) and

―B‖ (Temkin‘s constant related to heat of adsorption) were 0.03, and 2.6,

respectively. Regression coefficient (R2= 0.9997) was highly significant (Table

4.3). Linear plot of the linear relationship of ―qe‖ versus ―Ce‖ is presented in Fig.

4.17. ‖


―Pseudo-second order kinetics model, as previously described, was

employed for evaluation of removal of Cr (VI) ions by biosorption from dead

biomass of S. communis (the concentration of biomass was kept at 1 g/L while

sorbate concentrations were 100 mg/L, 200 mg/L, 300 mg/L, and 400 mg/L) to

explain the correlation between the equilibrium concentration of metal ions in

38

the solid phase (sorbent) and the aqueous solution. From the linear plot of ―t/q‖

versus ―t‖ (Fig. 4.18), the value of ―qe‖ was found to be 103.16 (slope of the

graph), the second order rate constant (K2) for the adsorption was calculated

as 0.0014 g mg-1 min-1 (derived from intercept of the plot), and the regression

coefficient (R2) was calculated as 0.9997 (Table 4.4).

‖

4.1.2.7. Thermodynamic studies

―Thermodynamic behavior of the adsorption of Cr (VI) on the surface of

S. communis was studied by calculating various thermodynamic constants. The

previously described equation in this context was used to obtain the values of

entropy (∆H) and enthalpy (∆S). The linear plot of the lnb Versus 1/T (Fig.

4.19) was drawn and by determining the intercept and slope, values of entropy

(∆H) was calculated as 4.32 J/mol and enthalpy (∆S) as 29.30 J mol-1K-1,

respectively (Table 4.5). ‖

―Arhenius equation (already described) used to obtain the values of

Arhenius constant (A₀) and the activation energy (Ea) gave the values of 52.1

and 13.46 J mol-1g-1, respectively (Table 4.6). Fig. 4.20 shows a linear plot of

―lnK2‖ versus ―1/T‖ (Arhenius equation). The values of ∆G for sorption of Cr (VI)

ions on S. communis were calculatedby using already given equation and the

values amounted to: 13.25, 11.71, 11.89, and 12.78 at 283 K°, 293 K°, 303 K°,

and 313 K°, respectively (Table 4.7). ‖

4.2. BIOSORPTION STUDIES ON C. delmatica

―4.2.1. Biosorption of Cadmium

―4.2.1.1. Effect of pH


4.0, 5.0, 6.0 and 7.0 (biosorbent concentration at 1g/L; initial sorbate


of 1.0, increasing gradually with the increase in pH values. The maximum

biosorption of 55 mg/g was achieved at pH 5.0, beyond this pH value (up to pH:

7.0) biosorption capacity declined (Fig. 4.21). ‖

39



importance as evident from Fig. 4.22 showing the biosorption efficiency of Cd

(II) by C. delmatica as a function of contact time (biosorbent: 1g/L; initial

sorbate concentration: 100 mg/L; pH: 5.0). Maximum biosorption took place in

the first 30 min., of contact time (50 mg/g), followed by slower uptake up to 90

min. and thereafter, equilibrium was probably achieved (qeq= 55 mg/g) because

no significant uptake of Cd (II) ions after 90 min., of exposure.‖


―Temperature is very important kinetic parameter in the biosorption





adsorption was maximum at 30oC (47 mg/g) followed by 20oC (Fig. 4.23), with

lower adsorption noticed at the temperatures of 10oC and 40oC. Temperature

also affects the time required to achieve maximum biosorption as was

observed in this study that the time required to achieve maximum biosorption

was reduced at elevated temperatures.‖






The biosorption of Cd (II) by C. delmatica was found to be directly proportional

to the biomass concentration. However, at exposure of low sorbate


gram biosorbent (qe) was maximum, achieving a value of 55 mg/g whereas at 3

g/L dose, it was only 26 mg/g (Fig. 4.24). ‖

40

4.2.1.5. Biosrption Equilibrium Isotherm

―This study was performed at pH value of 5.0, biosorbent concentration

of 1 g/L, and varying sorbate concentrations of 10, 20, 30, 40 mg/L.‖



concentrations of Cd (II) ions onto C. delmatica, equilibrium metal ion uptake



These values were then used to draw Langmuir isotherm according to the

equation already discussed.‖

―The linear plot of ―1/qe versus 1/Ce for biosorption of Cd (II) by C.

delmatica is shown in Fig. 4.25. A typical equilibrium biosorption isotherm is

obvious from the figure, suggesting that biosorption of Cd (II) ions involves a

chemical equilibrated and saturable mechanism which reflects site-specific

biosorption on the surface of the sorbent. ―qmax‖ was found to be 11.9 mg/g, the

value of ―b‖ was calculated as 0.06 Lmg-1, and ―R2‖ value as calculated as

0.9990. The value of ―k‖ for Cd (II) ions between the adsorbent and the

aqueous solution (qe/Ce) was 1.17.



Cd (II) towards C. delmatica surface by employing Freundlich equation, already

stated.

Fig. 4.26 shows a linear relationship of ―Log qe‖ versus ―Log Ce‖. The

value of Freundlich‘s saturation constant (KF) was found to be 7.72, constant

(n) was 0.48 (derived from intercept and slope respectively), and regression

coefficient value was highly significant (R2 = 0.9959). (Table 4.6) ‖


Temkin model was employed to the data derived from biosorption of Cd

(II) by C. delmatica, linear plot of qe versus Ce was drawn (Fig. 4.27) and the

values of ―KT‖ was calculated as 0.446, the constant, ―B‖ was 0.48 (derived

41

from the slope and the intercept values, respectively), and the regression

coefficient (R2 = 0.9999) very highly significant (Table 4.3). ‖



biomass of C. delmatica, was investigated (biomass concentration: 1 g/L,

pH=5.0, and initial sorbate concentrations: 10 mg/L). The data was employed to

the pseudo-second order equation already discussed (Ho and McKay, 1999).


was found to be 54.8 (slope of the graph), ―K2‖ for the adsorption mounted to

0.08 g mg-1 min-1 (derived from intercept of the plot), and the ―R2‖ (0.9998)

value was highly significant (Table 4.4).‖


―Thermodynamic behavior of the adsorption of Cd (II) ions on to the

surface of C. delmatica was studied by calculating various thermodynamics

constants. By applying the equation already discussed, ―∆H‖ and ―∆S‖ were

determined as 4.8 J/mol and 23.1 J mol-1K-1 (Table 4.5).The linear plot of the

lnb Versus 1/T is shown in Fig. 4.29.

― Arhenius equation was used to draw a linear plot of ―ln K2‖ versus

―1/T‖ (Fig. 4.30) and to obtain the values of Arhenius constant (A₀) and the

activation energy (Ea) as 1.5 and 35.58 J mol-1g-1, respectively (Table 4.6) .

―The values ―∆G‖ for sorption of Cd (II) ions on C. delmatica were

calculated as 13.24, 21.7, 26.2, and 16.48 at temperatures, 283 K°, 293 K°,

303 K°, and 313 K°, respectively (Table 4.7).

4.2.2. Biosorption of Chromium


―The biosorption of Cr (VI) was studied at varying pH values, such as

1.0, 2.0, 3.0, 4.0, 5.0, 6.0, and 7.0 pH units (initial sorbate concentration of 400

mg/L; sorbent concentration of 1g/L). The optimum adsorption of Cr (VI) was

recorded at a pH value of 4.0 (242 mg/g) (Fig.4.31). With the declining pH (> 4),

42

decrease in Cr (VI) biosorption was observed. Whereas, with the increase in

pH above 4.0 (up to pH: 7.0), decrease of biosorption was observed.


―Contact time of the exposure of biosorbent (C. delmatica) to the sorbate

[Cr (VI)] was observed as an important factor in biosorption application. Fig.

4.32 shows the biosorption efficiency of Cr (VI) by C. delmatica as a function of

contact time (initial sorbate concentration: 400 mg/L; sorbent concentration:

1g/L). Rapid uptake of Cr (VI) was observed in the first 30 min. (160 mg/g),

followed by relatively slower uptake up to 120 min. (242 mg/g); and thereafter,

no significant uptake of Cr (VI) ions was observed.





maximum at 20oC (247 mg/g), followed by 30oC (Fig. 4.33). However, lower







solution (Fig. 4.34). The biosorption capacity of Cr (VI) by C. delmatica was


1mg/L biosorbent concentration, maximum metal uptake per unit volume was

observed (242 mg/g) as evident from Fig. 4.34. ‖

43


―This study was performed at pH: 4.0; biosorbent concentration was 1

g/L, and different concentration of Cr (VI) ion solutions were 100, 200, 300, 400

mg/L.‖



using the equation described previously. The linear plot of‖1/qe‖ ―1/Ce‖ at

different initial Cr (VI) ions concentration in solution is shown in Fig. 4.35. This

figure shows a typical equilibrium biosorption isotherm, suggesting that

biosorption of Cr (VI) ions involves a chemical equilibrated and saturable

mechanism which reflects site-specific biosorption on the surface of the

sorbent. Maximum uptake capacity (Langmuir‘s qmax: 312 mg/g) of C. delmatica

was calculated; the value of saturation constant (b) was calculated as 0.04

L/mg, and the Regression coefficient was highly significant (R2 = 0.9142)

(Table 4.1). The value of ―k‖ (qe/ce) for Cr (VI) ions biosorption on to C.

delmatica was found to be 1.02 (Table 4.1).‖



estimate the adsorption intensity of the biosorbent biomass towards the sorbate

[Cr (VI)]; a linear relationship of ―Log qe‖ versus ―Log Ce‖ is shown in Fig. 4.36.

The values of KF was calculated to be 33.3, ―n‖ was 0.71 (derived from

intercept and slope respectively), and R2, 0.4542 (Table 4.2). ‖


―Temkin isotherm was employed as previosly described and the values

of ―KT‖ and ―B‖ were 0.03, and 2.6, respectively. Regression coefficient (R2=

0.9997) was highly significant (Table 4.3). Linear plot of the linear relationship

of ―qe‖ versus ―Ce‖ is presented in Fig. 4.37. ‖




44

biomass of C. delmatica (the concentration of biomass was kept at 1 g/L while

sorbate concentrations were 100 mg/L, 200 mg/L, 300 mg/L, and 400 mg/L) to

explain the correlation between the equilibrium concentration of metal ions in

the solid phase (sorbent) and the aqueous solution. From the linear plot of ―t/q‖

versus ―t‖ (Fig. 4.38), the value of ―qe‖ was found to be 54.8 (slope of the

graph), ―K2‖ was calculated as 0.08 g mg-1 min-1 (derived from intercept of the

plot), and ―R2‖ as 0.9941 (Table 4.4).‖


―Thermodynamic behavior of the biosorption of Cr (VI) on the surface of

C. delmatica was studied by calculating various thermodynamic constants. The

previously described equation in this context was used to obtain the values of

entropy (∆H) and enthalpy (∆S). The linear plot of the lnb Versus 1/T (Fig.

4.39) was drawn and by determining the intercept and slope, values of ―∆H‖

was calculated as 2.2.8 J/mol and ―∆S‖ as 52.29 J mol-1K-1, respectively (Table

4.5). ‖

―Arhenius equation (already described) was used to obtain the values of

―A₀” and ―Ea‖ gave the values of 17.6 and 10.39 J mol-1g-1, respectively (Table

4.6). Fig. 4.40 shows a linear plot of ―lnK2‖ versus ―1/T‖ (Arhenius equation).

The values of ∆G for sorption of Cr (VI) ions on C. delmatica were calculated by

using already given equation and the values amounted to: 12.94, 12.09, 12.31,

and 13.14 at 283 K°, 293 K°, 303 K°, and 313 K°, respectively (Table 4.7). ‖

4.3. STUDIES ON Spirogyra spp.

―4.3.1. Biosorption of Cadmium

―4.3.1.1. Effect of pH


4.0, 5.0, 6.0, and 7.0 (biosorbent concentration at 1g/L; initial sorbate


of 1.0, increasing gradually with the increase in pH values. The maximum

45

biosorption of 65 mg/g was achieved at pH 5.0, beyond this pH value (up to pH:

7.0) biosorption capacity declined (Fig. 4.41). ‖



importance as evident from Fig. 4.42 showing the biosorption efficiency of Cd

(II) by Spirogyra spp.as a function of contact time (biosorbent: 1g/L; initial

sorbate concentration: 100 mg/L; pH: 5.0). Maximum biosorption took place in

the first 30 min., of contact time (25 mg/g), followed by slower uptake up to 120

min. and thereafter, equilibrium was probably achieved (qeq= 47 mg/g) because

no significant uptake of Cd (II) ions after 120 min., of contact time was noticed.‖


―Temperature is a very important kinetic parameter in the biosorption





adsorption was maximum at 20oC (47 mg/g) followed by 30oC (Fig. 4.43), with

lower adsorption noticed at the temperatures of 10oC and 40oC. Temperature

also affects the time required to achieve maximum biosorption as was

observed in this study that the time required to achieve maximum biosorption

was reduced at elevated temperatures.‖






The biosorption of Cd (II) by Spirogyra spp., was found to be directly

proportional to the biomass concentration. However, at exposure of low sorbate


gram biosorbent (qe) was maximum, achieving a value of 47 mg/g whereas at 3

g/L dose, it was only 19 mg/g (Fig. 4.44). ‖

46

4.3.1.5. Biosrption Equilibrium Isotherm

―This study was performed at pH value of 5.0, biosorbent concentration

of 1 g/L, and varying sorbate concentrations of 10, 20, 30, 40 mg/L.‖



concentrations of Cd (II) ions onto Spirogyra spp., equilibrium metal ion uptake



These values were then used to draw Langmuir isotherm according to the

equation already discussed.‖

―The linear plot of ―1/qe versus 1/Ce for biosorption of Cd (II) by Spirogyra

spp., is shown in Fig. 4.45. A typical equilibrium biosorption isotherm is obvious

from the figure, suggesting that biosorption of Cd (II) ions involves a chemical

equilibrated and saturable mechanism which reflects site-specific biosorption

on the surface of the sorbent. ―qmax‖ was found to be 14.42 mg/g, the value of

―b‖ was calculated as 0.03 Lmg-1, and ―R2‖ value as calculated as 0.9883. The

value of ―k‖ for Cd (II) ions between the adsorbent and ―k‖ (qe/Ce) was 0.85. ‖



Cd (II) towards Spirogyra spp., surface by employing Freundlich equation,

already stated.

Fig. 4.46 shows a linear relationship of ―Log qe‖ versus ―Log Ce‖. The

value of Freundlich‘s saturation constant (KF) was found to be 2.46, constant

(n) was 0.39 (derived from intercept and slope respectively), and regression

coefficient value was highly significant (R2 = 0.9957). (Table 4.6) ‖


Temkin model was employed to the data derived from biosorption of Cd

(II) by Spirogyra spp.The linear plot of qe versus Ce was drawn (Fig. 4.47) and

the values of ―KT‖ was calculated as 0.781, the constant, ―B‖ was 0.01 (derived

from the slope and the intercept values, respectively), and the regression

coefficient (R2 = 0.9999) very highly significant (Table 4.3). ‖

47



biomass of Spirogyra spp., was investigated (biomass concentration: 1 g/L,

pH=5.0, and initial sorbate concentrations: 10 mg/L). The data was employed to

the pseudo-second order equation already discussed (Ho and McKay, 1999).


was found to be 57.54 Cal (slope of the graph), ―K2‖ for the adsorption was 0.04

g mg-1 min-1 (derived from intercept of the plot), and the ―R2‖ (0.9998) value

was highly significant.‖


―Thermodynamic behavior of the adsorption of Cd (II) ions on to the

surface of Spirogyra spp., was studied by calculating various thermodynamics

constants. By applying the equation concerning enthalpy and entropy already

discussed, ―∆H‖ and ―∆S‖ were determined as 2.2 J/mol and 70.7 J mol-1K-1

(Table 4.5). The linear plot of the lnb versus 1/T is shown in Fig. 4.49.

― Arhenius equation (as given earlier) was used to draw a linear plot of

―ln K2‖ versus ―1/T‖ (Fig. 4.50) and to obtain the values of Arhenius constant

(A₀) and the activation energy (Ea) as 1.23 and 38.34 J mol-1g-1, respectively

(Table 4.6) .

―The values of ―∆G‖ for sorption of Cd (II) ions on Spirogyra spp., were

calculated as 12.70, 11.8, 11.73, and 13.06 at temperatures of 283 K°, 293 K°,

303 K°, and 313 K°, respectively (Table 4.7).

4.3.2. Biosorption of Chromium


―The biosorption of Cr (VI) by Spirogyra spp. was studied at varying pH

values, such as 1.0, 2.0, 3.0, 4.0, 5.0, 6.0, and 7.0 pH units (initial sorbate

concentration of 400 mg/L; sorbent concentration of 1g/L). The optimum

adsorption of Cr (VI) was recorded at a pH value of 4.0 (265 mg/g) (Fig.4.51).

With the declining pH (> 4), decrease in Cr (VI) biosorption was observed.

Whereas, with the increase in pH above 4.0 (up to pH: 7.0), decrease of

biosorption was observed.

48


―Contact time of the exposure of biosorbent (Spirogyra spp.) to the

sorbate [Cr (VI)] was observed as an important factor in biosorption application.

Fig. 4.52 shows the biosorption efficiency of Cr (VI) by Spirogyra spp.as a

function of contact time (initial sorbate concentration: 400 mg/L; sorbent

concentration: 1g/L). Rapid uptake of Cr (VI) was observed in the first 30 min.

(220 mg/g), followed by relatively slower uptake up to 120 min. (272 mg/g); and

thereafter, no significant uptake of Cr (VI) ions was observed.





maximum at 20oC (252 mg/g), and at 30oC as well (Fig. 4.53). However, lower







solution (Fig. 4.54). The biosorption capacity of Cr (VI) by Spirogyra spp., was


1mg/L biosorbent concentration, maximum metal uptake per unit volume was

observed (272 mg/g) as evident from Fig. 4.54. ‖


―This study was performed at pH: 4.0; biosorbent concentration was 1

g/L, and different concentration of Cr (VI) ion solutions were 100, 200, 300, 400

mg/L.‖



using the equation described previously. The linear plot of‖1/qe‖ ―1/Ce‖ at

different initial Cr (VI) ions concentration in solution is shown in Fig. 4.55. This

49

figure shows a typical equilibrium biosorption isotherm, suggesting that

biosorption of Cr (VI) ions involves a chemical equilibrated and saturable

mechanism which reflects site-specific biosorption on the surface of the

sorbent. Maximum uptake capacity (Langmuir‘s qmax: 498 mg/g) of Spirogyra

spp., was calculated; the value of ―b‖ was calculated as 0.008 L/mg, and the

Regression coefficient was significant (R2 = 0.9012) (Table 4.1). The value of

―k‖ (qe/ce) for Cr (VI) ions biosorption on to Spirogyra spp., was found to be 0.88

(Table 4.1).‖



estimate the adsorption intensity of the biosorbent biomass towards the sorbate

[Cr (VI)]; a linear relationship of ―Log qe‖ versus ―Log Ce‖ is shown in Fig. 4.56.

The values of KF was calculated to be 108.76, ―n‖ was 0.31 (derived from

intercept and slope respectively), and R2, 0.6712 (Table 4.2). ‖


―Temkin isotherm was employed as previosly described and the values

of ―KT‖ and ―B‖ were 10.16, and 3.81, respectively. Regression coefficient (R2=

0.9999) was highly significant (Table 4.3). Linear plot of the linear relationship

of ―qe‖ versus ―Ce‖ is presented in Fig. 4.57. ‖




biomass of Spirogyra spp., (the concentration of biomass was kept at 1 g/L

while sorbate concentrations were 100 mg/L, 200 mg/L, 300 mg/L, and 400

mg/L) to explain the correlation between the equilibrium concentration of metal

ions in the solid phase (sorbent) and the aqueous solution. From the linear plot

of ―t/q‖ versus ―t‖ (Fig. 4.58), the value of ―qe‖ was found to be 98.92 Cal (slope

of the graph), ―K2‖ was calculated as 0.0009 g mg-1 min-1 (derived from

intercept of the plot), and R2 as 0.9959 (Table 4.4).‖

50


―Thermodynamic behavior of the biosorption of Cr (VI) on the surface of

Spirogyra spp., was studied by calculating various thermodynamic constants.

The previously described equation in this context was used to obtain the values

of entropy (∆H) and enthalpy (∆S). The linear plot of the lnb Versus 1/T (Fig.

4.59) was drawn and by determining the intercept and slope, values of ―∆H‖

was calculated as 2.73 J/mol and ―∆S‖ as 52.77 J mol-1K-1, respectively (Table

4.5). ‖

―Arhenius equation (already described) was used to obtain the values of

―A₀” and ―Ea‖ gave the values of 18.09 and 12.76 J mol-1g-1, respectively (Table

4.6). Fig. 4.60 shows a linear plot of ―ln K2‖ versus ―1/T‖ (Arhenius equation).

The values of ∆G for sorption of Cr (VI) ions on Spirogyra spp., were calculated

by using already given equation and the values amounted to: 12.7, 11.8, 11.73,

and 13.06 at 283 K°, 293 K°, 303 K°, and 313 K°, respectively (Table 4.7). ‖

4.4. STUDY OF BIOSORPTIVE BINDING SITES ON ALGAL

SURFACE

―Qualitative studies regarding binding sites of heavy metals on the

surface of biosorbents under investigation (S. communis, C. delmatica,

Spirogyra spp.) were carried out by Fourier Transform Infra-Red (FTIR)

Spectroscopy and Scanning Electron Microscopy (SEM).‖

4.4.1. FTIR Spectroscopy

4.4.1.2. Studies with S. communis

―FTIR spectra of S. communis showed a shift in peaks of after Cd (II)

ions treatment (Fig. 4.64.a,b). Shifts of peaks were observed as: 3369 cm-1

(unloaded samples) to 3376 cm-1 [Cd (II)-loaded samples]; 2944 cm-1

(unloaded samples) to 2938 cm-1 [Cd (II)-loaded samples]; and 1032 cm-1

(unloaded samples) to 1042 cm-1 [Cd (II)-loaded samples]. Whereas the above

mentioned peaks represent hydroxyl bonds, alkane stretches, and amine

bonds, of various functional groups, such as aldehyde, ketone, amine,

51

carboxylic acids, alcohols, ethers, esters, etc. All of these functional groups are

electronegative and are capable of attracting cations, such as Cd (II) (Table

4.8). ‖

―FTIR spectra of S. communis showed a shift in peaks of after Cr (VI)

ions treatment (Fig. 4.64.a,c). Shifts of peaks were observed as: 3369 cm-1

(unloaded samples) to 3312 cm-1 [Cr (VI)-loaded samples]; 2944 cm-1

(unloaded samples) to 2951 cm-1 [Cr (VI)-loaded samples]; and 1032 cm-1

(unloaded samples) to 1194 cm-1 [Cr (VI)-loaded samples]. Whereas the above

mentioned peaks represent carbonyl bonds of various functional groups, such

as aldehyde, ketone, carboxylic acids, alcohols, ethers, esters, etc. All of these

functional groups are electronegative and are capable of attracting cations,

such as Cr (VI) ion (Table 4.8).‖

4.4.1.3. Studies with C. delmatica

―FTIR spectra of C. delmatica showed a shift in peaks of after Cd (II)


(unloaded samples) to 3337 cm-1 (loaded samples); 2040 cm-1 (unloaded

samples) to 2022 cm-1 (loaded samples); and 1290 cm-1 (unloaded samples) to

1300 cm-1 (loaded samples). Whereas the above mentioned peaks represent

hydroxyl bonds, alkane stretches, and amine bonds, of various functional

groups, such as aldehyde, ketone, amine, carboxylic acids, alcohols, ethers,

esters, etc. All of these functional groups are electronegative and are capable

of attracting cations, such as Cd (II) (Table 4.8). ‖

―FTIR spectra of C. delmatica showed a shift in peaks of after Cd (II)


(unloaded samples) to 1522 cm-1 (loaded samples); and 1050 cm-1 (unloaded

samples) to 1060. Whereas the above mentioned peaks represent carbonyl

bonds of various functional groups, such as aldehyde, ketone, carboxylic acids,

alcohols, ethers, esters, etc. All of these functional groups are electronegative

and are capable of attracting cations, such as Cr (VI) ion (Table 4.8).‖

52

4.4.1.3. Studies with Spirogyra spp.

―FTIR spectra of Spirogyra spp. showed a shift in peaks of after Cd (II)


(unloaded samples) to 3337 cm-1 (loaded samples); 2040 cm-1 (unloaded

samples) to 2022 cm-1 (loaded samples); and 1290 cm-1 (unloaded samples) to

1300 cm-1 (loaded samples). Whereas the above mentioned peaks represent

hydroxyl bonds, alkane stretches, and amine bonds, of various functional

groups, such as aldehyde, ketone, amine, carboxylic acids, alcohols, ethers,

esters, etc. All of these functional groups are electronegative and are capable

of attracting cations, such as Cd (II) (Table 4.8). ‖

―FTIR spectra of Spirogyra spp. showed a shift in peaks of after Cd (II)


(unloaded samples) to 1522 cm-1 (loaded samples); and 1050 cm-1 (unloaded

samples) to 1060. Whereas the above mentioned peaks represent carbonyl

bonds of various functional groups, such as aldehyde, ketone, carboxylic acids,

alcohols, ethers, esters, etc. All of these functional groups are electronegative

and are capable of attracting cations, such as Cr (VI) ion (Table 4.8).‖

4.4.2. Scanning Electron Microscopic (SEM) Studies

Scanning electron microscopy (SEM) were performed for the dead

biomass of all the three algae under investigation (S. communis, C. delmatica,

Spirogyra spp.), to elucidate the binding of Cd (II) and Cr (VI) ions on the algal

surface. Also, SEM facilitated observation at higher magnification of closely

spaced materials, such as binding sites, etc., on the algal cell wall.

Cd (II) ions loaded [Biosorbents under investigation exposed to 100mg/L

of Cd (II) ion solution] and Cr (VI) ions loaded [Biosorbents under investigation

exposed to 100mg/L of Cr (VI) ion solution] were subjected to SEM studies.

Also, biosorbents exposed to pure distilled water (without sorbate) were studied

with SEM (control).‖

SEM studies revealed the topographical and elemental appearance of

the algal samples under investigation under a virtually large depth of field,

allowing different specimen parts to stay in focus at a time.

53

―Cd (II) loaded biosrbents [exposed to 100 mg/L Cd (II) ion solution at

pH 4 (optimum conditions for biomass of Cd (II) ions by S. communis, C.

delmatica and Spirogyra spp.)], showed a difference in morphology to the

unloaded samples. In the case of loaded samples, the matrix layers of the cell

wall were observed to shrink and stick (Fig. 4.61-a, 4.61-b, 4.61-b; Fig. 4.62-a,

4.62-b, 4.62-c; 4.63-a, 4.63-b, 4.63-c). In case of Cr (VI) biosorption, the

topological change after biosorption was even more pronounced.‖

54

Table 3.1. Composition of Bold 1NV medium (UTEX, the culture

collection of algae, USA)

Component Amount Stock Solution Concentration

Final Concentration

1 NaNO3 (Fisher BP360-500)

10 mL/L 10 g/400mL dH2O 2.94 mM

2 CaCl2·2H2O (Sigma C-3881)

10 mL/L 1 g/400mL dH2O 0.17 mM

3 MgSO4·7H2O (Sigma 230391)

10 mL/L 3 g/400mL dH2O 0.3 mM

4 K2HPO4 (Sigma P 3786)

10 mL/L 3 g/400mL dH2O 0.43 mM

5 KH2PO4 (Sigma P 0662)

10 mL/L 7 g/400mL dH2O 1.29 mM

6 NaCl (Fisher S271-500)

10 mL/L 1 g/400mL dH2O 0.43 mM

7 P-IV Metal Solution

6 mL/L

8 Vitamin B12 1 mL/L

9 Biotin Vitamin Solution

1 mL/L

10 Thiamine Vitamin Solution

1 mL/L

http://www.sbs.utexas.edu/utex/mediaDetail.aspx?mediaID=127







55

Table 4.1. Langmuir Isotherms’ constants for the sorption of Cd

(II) and Cr (VI) ions by the investigated algal

biosorbents

Cd Cr

Biosorbent qmax

(mg/g

)

b

(L/mg)

R2

k qmax

(mg/g)

b

(L/mg)

R2 K

S. cummunis 1.44 2.5 .960

5

0.7

5

411 .01 .9998 1.08

C. delmatica 11.9 0.06 .999

0

1.17 312 .04 .9142 1.02

Spirogyra

spp.

14.42 0.03 .988

3

0.85 498 .008 .9012 0.88

56

Table 4.2. Freundlich Isortherms’ constants for the sorption of Cd (II)

and Cr (VI) ions by the investigated algal biosorbents

Cd Cr

Biosorbent n KF R2 n KF R2

S. cummunis 0.16 14.49 .9986 0.85 27.38 0.9988

C. delmatica 0.48 7.72 .9959 0.71 33.3 0.4542

Spirogyra spp. 0.39 2.46 .9957 0.31 108.76 0.6712

57

Table 4.3. Temkin Isortherms’ constants for the sorption of Cd (II) and

Cr (VI) ions by test biosrbents

Cd Cr

Biosorbent KT B R2 KT B R2

S. cummunis 0.024 4.31 0.9999 0.03 2.69 0.9997

C. delmatrics 0.446 0.48 0.9997 0.1 0.20 1

Spirogyra spp. 0.781 0.01 1 10.16 3.81 0.9999

58

Table 4.4. Pseudo-second order kinetic constants for the biosorption of

Cd (II) and Cr (VI) ions by the investigated algal biomass

Biosorbent Cd (II) Cr (VI)

q eq Cal K2 R2 q eq Cal K2 R2

S. cummunis 43.61 0.1 0.9999 103.16 0.0014 0.9997

C. delmatica 54.80 0.08 0.9998 108.26 0.0012 0.9994

Spirogyra spp. 57.54 0.04 0.9941 98.92 0.0009 0.9959

59

Table 4.5. Thermodynamic parameters, enthalpy and entropy for the

biosorption of Cd (II) and Cr (VI) ions by the investigated alga

Cd (II) Cr (VI)

Biosorbent ∆H°

(KJmol-1)

∆S°

(J mol-1K-1)

∆H°

(KJmol-1)

∆S°

(J mol-1K-1)

S. cummunis 8.5 41.8 4.32 29.30

C. delmatica 4.8 23.1 2.28 52.29

Spirogyra spp. 2.2 70.7 2.73 52.77

60

Table 4.6. Thermodynamic parameters, Arhenius constant (A₀) and

Energy of activation (Ea) for the biosorption of Cd (II) and Cr

(VI) ions by the investigated algal biosorbents

Cd (II) Cr (VI)

Biosorbent A₀ Ea (J mol-1g-

1)

A₀ Ea(J mol-1g-1)

S. cummunis 4.2 35.58 52.1

13.46

C. delmatica 1.5

38.17 17.6

10.39

Spirogyra spp.

1.23

38.34 18.09

12.76

61

Table 4.7. -∆G°(KJmol-1) values for biosorption of Cr (VI) and Cd (II)

ions by the investigated algal biosorbent at varying

temperatures

Temperature (Ko): 283 293 303 313

Cd (II)

S. cummunis 12.31 20.03 21.95 11.43

C. delmatica 13.24 21.70 26.20 16.48

Spirogyra spp. 14.89 24.40 24.35 12.07

Cr (VI)

S. cummunis 13.25 11.71 11.89 12.78

C. delmatica 12.94 12.09 12.31 13.14

Spirogyra spp. 12.70 11.80 11.73 13.06

62

Table 4.8. FTIR absorption wavelengths (Cm-1 ) for Cd (II) ion and Cr (IV)

ions biosorption by the investigated algal biomass

Cd

Treatme

nt

Function

al group

Cr

Treatme

nt

Function

al group

Unloaded Loade

d

Unloaded loade

d

S.cummun

is

3369 3376 −OH,

−NH

3369 3312 −C=O

2944 2938 −CH 2944 2951 −C−O

C.delmatic

a

2921 2905 −OH 1032 1194 −C=O

3288 3296 −CH 3288 3266 −C−O

1637 1631 −C-H 1030 1020 −C=O

Spirogyra

spp.

3354

2921

3337

2905

−OH,

−NH

−OH

3354

1650

3352

1634

−OH,

−NH

−C=O

1650 1634 −CH 1650 1641 −CH

1538 1557 −C-H 1020 1032 −C−O

63

T

I

M

E

73

‖

76

Biomass 1 g

123

Figure 4.61 (a). SEM of S. communis unloaded (CONTROL).

Figure 4.61. (b) SEM of S. communis , Cd (II)-loaded

Figure 4.61 (c). SEM of S. communis , Cr (VI)-loaded

124

(a)

(b)

(c ) Figure 4.62 (a) SEM of C. delmatica unloaded (CONTROL)

Figure 4.62 (b) SEM of C. delmatica Cd loaded

Figure 4.62 (c) SEM of C. delmatica Cr loaded

125

(a)

(b)

(c) Figure 4.63. (a) SEM of Spirogyra spp. (unloaded control)

(b) SEM of Spirogyra spp. [Cd (II)-loaded]

(c) SEM of Spirogyra spp. [Cr (VI)-loaded]

126

(a)

(b)

© 1000

4000

00

1000 4000

1000 4000

Wave number cm-1

Wave number cm-1

Wave number cm-1

Fig: 4.64. FTIR spectra of S. communis (a) Control (unloaded) (b) Cd (II) loaded (c) Cr (VI) loaded

127

(a)

(b)

( c)

500 4000

500 4000

500 4000

Fig: 4.65. FTIR spectra of C. delmatica (a) Control (unloaded) (b) Cd (II) ions loaded (c) Cr (VI) ions loaded

Wave number cm-1

Wave number cm-1

Wave number cm-1

128

(a)

(b)

(c)

Fig: 4.64. FTIR spectra of Spirogyra spp. (a) Control (unloaded) (b) Cd (II) ions loaded (c) Cr (VI) ions loaded

500

500

500

4000

4000

4000 Wave number cm-1

Wave number cm-1

Wave number cm-1

129

O

OH

OH

OH

R1OH

R1

OH O

OH

O-

O-

R1OH

R1

OH2H2O(1)

-2H2O(2)

+ 2H3O+(aq)

O

OH

OH

OH

R1OH

R1

OH O

OH

O-

O-

R1OH

R1

OH

M2+

2H2O(1)

-2H2O(2)

+ 2H3O+(aq)+ M

2+(aq)

a

b

Figure 5.1 Involvement of functional groups in providing divalent

metal ions (M2+) a site of attachment

(a) Available functional groups at optimum pH

(b) Binding of metal ions on the electronegative functional

groups (available binding sites)

130

O H

O

Cr

CH3

O H

O

OH2

OH2

O H

O

Cr

CH3

O H

O

Fig 5.2 Incorporation of Cr (VI) ion in the organic components on the

cell wall of green algae.

131

R C OH

O

+ OH2 R C O-

O

+ H3O+

R C

O_

O

R C

O

O

R C

O

O

Carboxylate ion

Figure 5.3 The carboxyl ion on the surface of algae provide

site of biosorption onto the surface of biosorbets, studied in

the present investigation.

132

P O

O

O

O

Ca2+

P O

O

O

O

Cd2+

Cd2+

Ca2+

(A)

(B) OH

O

OH

O

O OO

RO OH

CH2OHOH

CH2 OH

O

O

Cd

O

OCH3OCH3

O

NH2

NH2

II

III

I

Figure: 5.4 The mechanism of ion exchange at electonegative

functional group:

(A) Ca2+ is exchanged by Cd+2 at phosphate group

(B) Cd 2+ ion may be incorporated by coordinate covalent

links

133

O

CH2OH

OH

O

OO

CH2OH

OH

OH

OH

CH2OH

OH

CH2OH

OHOH

OH

O

O

O

O

CH2OH

OH

O

OO

CH2OH

OH

OH

OH

CH2OH

OHOH

OH

O

O

O

OH

H,OH

H,OH

80

5

OH

CH2OH

O

O

OH

OH

OH

OH

- - - - - - - - O

- - - - - - - - O

O

O

OH

O

O

O

O

O

O

OH

(1 4)-linked

(1 3)-linked

(a)

(b)

(c)

Fig. 5.5. The fibrillar molecules of algal cell walls of green algae:

(a) Algalcellulose, a 𝛃 (1→4) linked unbranched glucan, of the

brownalgae;

(b) Structural units present in xylan from the red algae,both 𝛃 (1→3)

and 𝛃

(1→4) linked forms have been isolated;

(c) Mannan, a 𝛃 (1→4) linked d-mannose from the red algae. After

Percival

and McDowell (1967) .

134

CHAPTER 5

DISCUSSION

―Investigation of factors that affect the efficiency of biosorption of metal

ions of both Cd (II) and Cr (VI), by using dried biomass of filamentous green

algae taken from three different sources, Spirogyra communis, Cladophora

delmatica, and Spirogyra spp. is of vital importance in terms of understanding

the mechanism(s) of biosorption process. The use of analytical techniques,

such as SEM and FTIR have facilitated in the development of comprehensive

understanding of the binding sites for biosorption of heavy metal(s) by the algal

species employed. This knowledge shall also be of great interest to the

industrial outfits, such as wastewater treatment and its engineering, enabling

them for technological enhancement.‖

5.1 BIOSORPTION CAPACITY OF THE BIOSORBENTS

―In the present study, some of the critically important physico-chemical

characteristics, such as pH, contact time, temperature, biosorbent

concentration, etc., affecting the biosorption capacity of Cd (II) and Cr (VI) ions

by S. communis, C. delmatica, and Spirogyra spp. were studied under varying

pH, contact time, biosorbent concentration, and temperature.‖

5.1.1. Effect of pH

――The pH of the solution seems to have a significant influence on

dissociation of the surface of algal biomass and the solution chemistry of the

heavy metals; for example, hydrolysis, complexation by organic and/or

inorganic ligands, redox reactions, and precipitation, as well as the speciation

and biosorption availability for heavy metals. Similar previous studies on

biosorption of heavy metals by using dead biomass of various organisms

strengthens this evidence (Arief, et al., 2008; Wang and Chen, 2006; Pavasant

et al., 2006; Guo et al., 2008; Komy et al., 2006; Sheng et al., 2004; Romera et

al., 2006). Hence, the effect of pH on metal uptake can be associated with both

135

the surface functional groups on the biomass' cell walls as well as on the metal

chemistry of the solution (Sheng et al., 2004). Equilibrium of the system is also

affected by pH (Romera et al., 2006). The Following equations explain the

dependence of biosorption on pH of the solution (Arief et al., 2008). ‖

―Where [AH] and [A-] are concentrations of protonated and deprotonated

surface functional groups, respectively. When the pH value is lower than the

pKa, the equilibrium shifts to the left (as shown below).‖

Protons are consumed and the pH rises until its value approaches the

pKa. The opposite trend happens if the pH value is higher than the pKa

(Romera et al., 2006).

―In the present study, Cd (II) ion uptake seems to be optimum at pH

value of 5.0 in case of all the three algal species (S. communis, C. delmatica,

and Spirogyra spp.). In case of S. communis, uptake of Cd (II) ion was reduced

from 43 to 20%, when pH value was reduced from 5.0 to 1.0. The improved

metal sorption with higher pH was attributable to the increase in the amount of

ligands for metal ion binding. Moreover, at low pH, competition occurs between

protons and metal ions, leading to less metal uptake. By mechanism, at low pH,

the concentration of H+ ions was high. Hence Cd (II) ions must compete with H+

ions in order to attach to the surface functional groups of the olive cake. With

the rise in pH, fewer H+ ions exist, and consequently, Cd (II) ions have a better

chance to bind to the binding sites that are free. Later, when the pH enters

basic conditions, the formation of Cd(OH)3 takes place due to the dissolution of

Cd(OH)2 and as a result, the adsorption rate decreases as shown below:

136

――While increasing pH value from 5.0 to 7.0, a decrease from 43 to 24%

in Cd (II) ion uptake was recorded. The probable reason for this could be that at

higher pH of Cd (II) ions precipitate to Cd (OH)2 as indicated by an above

equation. Several previous studies have described similar trends as observed

in the present studies. Investigation of Pb (II), Cu (II), Cd (II), Zn (II), and Ni (II)

sorption by Sargassum sp. and Padina sp. (brown macroalgae), Ulva sp. (a

green macroalga), and Gracillaria sp. (a red macroalga) was conducted in the

pH range of 2.0 – 6.0 (Salem and Allia, 2008). Cd(II) by Corallina officinalis,

Porphyra columbina and Codium fragile (Cookierman, 2007); Cu(II) and Pb(II)

by Cladophora fascicularis (Deng et al., 2007); Pb(II) by Spirogyra sp. (Gupta

and Rastogi, 2008); Cr(III) by Chlorella miniata (Han et al., 2006); Cu(II), Cd(II),

Pb(II), and Zn(II) by Caulerpa lentillifera (Pavasant et al., 2006); Cd(II), Cu(II),

Ni(II), Pb(II) and Zn(II) by Codium vermilara, Spirogyra insignis, Asparagopsis

armata, Chondrus crispus, Fucus spiralis and Ascophyllum nodosum (Romera

et al., 2006) and also Pb(II) and Cd(II) by Ulta lactuca (Sari and Tuzen, 2008).

Al-Anber and Matouq (2008) looked at the pH effect on metal uptake in the

adsorption process of heavy metals using agricultural waste, such as olive

cake. Upon increasing the pH from 2 to 6, the removal efficiency of Cd (II) was

reported to increase from 42 to 66%. Pino et al. (2006) showed that the

removal of cadmium ion by green coconut shell powder was pH dependent.‖ ‖

――Chromium (VI) ion uptake by the biosorbent investigated in the present

study was found to be optimum at pH 4.0. With a decrease in pH from 4.0 to

1.0, Cr (VI), uptake decreased from 50 to 19% in S. communis, from 48 to 18%

in C. delmatica, and from 53 to 18% in Spirogyra spp. (Fig.4.1, 4.21, and 4.41).

This decreased uptake might be attributed to the fact that low pH offers

competition between protons and metal ions in the solution, leading to reduced

adsorption. Similar tendency was also observed in adsorption of Cr (VI) by

green algae (Bishnoi et al., 2007). The hexavalent chromium [Cr (VI)] ions exist

in the form of HCrO4 and Cr2O7 so that their equilibrium relations can be

expressed in the following two equations:‖ ‖

137

However, in alkaline pH, the equilibrium relation differs as represented by the

following equations:

― ―Hence, high adsorption capacity at low pH values (below pH 4.0) is

due to the presence of excess H+ ions that are able to neutralize the negatively

charged adsorbent surface, and thereby, reduce the hindrance for diffusion of

dichromate ions (Kumari et al., 2006). Decreased uptake of Cr (VI) ion was

observed owing to protonation of certain binding sites at algal biomass surface.

‖

――On the contrary, with the increase in pH from 4.0 to 7.0, Cr (VI) ion

uptake has been found to decrease from 50 to 28% in S. communis, from 48 to

19% in C. delmatica, and from 42% to 32% in Spirogyra spp., as observed in

the present study. At higher pH, elevated metal (at pH 4.0) sorption was

attributable to the increase in the number of available ligands (binding sites) for

metal ion binding, whereas the lower percentage of sorption of Cr (VI) ion onto

sorbent at elevated pH (5.0, 6.0, and 7.0) might be attributed to precipitation

reaction which occurs at pH 5.0-6.0 (Matheickal et al., 1999).‖ ‖

―Other studies regarding chromium ions adsorption on the surface of

various biomasses also seem to agree with the findings in the present study

(Ahalya et al., 2005; Mohanty et al., 2006; Malkoc and Nohuglu. 2007; Baral et

al., 2007; Das et al., 2007; Isa et al., 2008; Garg et al., 2007; Park et al., 2005;

Pino et al., 2006; Anjana et al., 2007; Hassan et al., 2007). ‖

5.1.2. Effect of Contact Time

―In the present study equilibrium of the solution was generally achieved

in 120 min. by S. communis, Cldophora delmatica, and Spirogyra spp. for the

sorption of Cd (II) ions and Cr (VI) ions. Initially (during the first 30 min.), rapid

uptake of cations from biosorption by algal biomass was recorded, slowing

138

down significantly after the first 90 min. This was probably because initially a

large number of vacant surface sites might have been available for adsorption

and after some time, the remaining vacant surface sites may have been difficult

to be occupied due to forces between the solute molecules of the solid and bulk

phase. Also, the diminishing removal of metal ions with increasing time may

have been due to intraparticle diffusion process dominating over adsorption

(Volesky, 2003; Deo and Ali, 1993). ‖

―The effect of contact time on the extent of adsorption of heavy metals

has been extensively recorded in numerous studies conducted on the

biosorption of heavy metal(s) by various organic sorbents and variations in time

required to achieve equilibrium in solution was found to be dependent upon the

sorbent. Equilibrium time for sorption of Cd (II) ion by brown algae has been

found to be 30 min. (Liu et al., 2007), by brown algae as 60 min. (Sari and

Tuzen, 2008), by S. condensate as 120 min. (Onyancha et al., 2008), by

Spirogyra neglecta and Cladophora calliceima as 50 min. (singh et al., 2006). ‖

5.1.3. Effect of Temperature

―The temperature has been found to affect the biosorption. The extent of

Cd (II) ion sorption by S. communis has been found maximum at 30oC (48%)

followed by its capacity at 20oC (46%), whereas low cation uptake has been

recorded at 10oC (30%) and 40oC (34%); by C. delmatica, maximum at 30oC

(47%) followed by its capacity at 20oC (43%), whereas low cation uptake was

recorded at 10oC (27%) and 40oC (31%); and by Spirogyra spp., maximum at

20oC (46%), followed by its capacity at 30oC (44%), whereas low cation uptake

has been recorded at 10oC (30%) and 40oC (34%). Similarly, the extent of Cr

(VI) ion sorption by S. communis has been found maximum at 20oC (52%)

followed by its capacity at 30oC (48%), whereas low cation uptake was

recorded at 10oC (30%) and 40oC (34%); by C. delmatica, maximum at 20oC

(49%) followed by its capacity at 30oC (45%), whereas low cation uptake has

been recorded at 10oC (32%) and 40oC (29%); and by Spirogyra spp.,

maximum at 30oC and 20oC (51%), whereas low cation uptake was recorded at

10oC (40%) and 40oC (32%). The decreased adsorption at lower temperature

can be due to shrinking of the surface and that at the elevated temperatures

139

can be attributed to exothermic characteristics of the biosorbent. The

temperature change alters the adsorption equilibrium in a specific way

determined by the exothermic or endothermic nature of a process. Previously,

a number of studies concerning the effect of temperature on isotherms, metal

uptake and also with respect to biosorption thermodynamics parameters have

been performed (Dundar et al., 2008; Benguella and Benaissa, 2002; Malkoc

and Nuhoglu, 2007; Gupta and Rastogi, 2008; Park et al., 2005; Hassan et al.,

2007; Ho, 2006). The impact of temperature on the adsorption isotherm of

metal ions by sorbent at a certain pH was explained by Shen and Duvnjak

(2004) in which higher temperatures were found to support the uptake of

cations. A more enhanced level of uptake in parallel with a temperature rise

resembles the nature of a chemisorption mechanism (endothermic process).

These trends were supported by other studies as well (Park et al., 2005; and

Guzel et al., 2008), except the one in which baker‘s yeast was found to show

better uptake of cations at lower temperatures (Padmavathy, 2008). ‖

5.1.4. Biosorbent concentration

―The biosorption capacity of sorbents was found to be directly

proportional to the biomass concentration (Aydin et al., 2008; Dundar et al.,

2008). However, the amount adsorbed per unit mass decreases. In the present

study, all the biosorbents, S. communis, C. delmatica, and Spirogyra spp.

showed better metal ion uptake (q values) at 1 g/L concentration as compared

to 2 g/L and 3g/L, per unit mass. This trend could be explained as a

consequence of a partial aggregation of biomass at higher biomass

concentration, which results in a decrease in effective surface area for the

biosorption (Gupta et al., 2001). The amount of biomass in the solution also

affects the specific metal uptake. In principle, with more adsorbent present, the

available adsorption sites or functional groups also increase (Gupta and

Rastogi, 2008; Tunali et al., 2006). In turn, the amount of adsorbed heavy metal

ions is increased, which brought about an improved adsorption efficiency

(Bueno et al., 2008; Mohanty et al., 2006; Baral et al., 2007; Popuri et al., 2007;

Garg et al., 2007; Isa et al., 2008, Akar et al., 2006; Romera et al., 2006; Al-

Anber and Matouq, 2008; Ghodbane et al., 2008; Bishnoi et al., 2007;Sari and

Tuzen, 2008; Ho and Ofomaja, 2006; and Cho and Kim, 2003). ‖

140

5.1.5. Biosorption Equilibrium Isotherms

―The analyses of equilibrium data are important for developing an

equation that can be used for design purposes. The Langmuir, Freundlich, and

Temkin adsorption constants and their correlation coefficients were evaluated

from the respective isotherms. Figures 4.5-4.7, 4.15-4.17, 4.25-4.27, 4.35-4.37,

4.45-4.47, and 4.55-4.57 show that the biosorption data of the present study

could be linearized to fit in various adsorption isotherms, such as Langmuir,

Freundlich and Temkin equations and the values of correlation coefficient for

the Langmuir and Temkin adsorption isotherms were found higher than those

found in Freundlich isotherm(Table 4.1, 4.2, and 4.3) for both the tested

sorbates [Cd(II) and Cr (VI)] by the sorbents under investigation (S. communis,

C. delmatica, and Spirogyra spp.). According to the results, the equilibrium

adsorption data was better fitted to Langmuir and Temkin adsorption isotherm

models than Freundlich adsorption model and generally high regression

correlation coefficients (>0.995) were found, suggesting that the Langmuir and

Temkin models are very suitable for describing the biosorption equilibrium of

Cd (II) and Cr (VI) by the S. communis, C. delmatica, and Spirogya spp. The

applicability of the Langmuir isotherm model to the metal ion-dried algae

system implies that monolayer biosorption conditions exist under the

experimental conditions used (Pahlvanzadeh et al., 2010). This also defines

that the binding energy on the whole surface of biosorbent was uniform.

Moreover, it can be inferred that the adsorbed metal ions do not interact or

compete with each other thus forming a monolayer. ‖

―The maximum capacity qmax determined from the Langmuir isotherm

defines the total capacity of the biosorbent for Cr (VI) and Cd (II) ions. The

adsorption capacity of biosorbent (qmax), increases with increase in temperature

(Figure 4.3, 4.13, 4.23, 4.33, 4.43, 4.53). The present study seems to agree

with the trend shown by similar biosorption studies previously conducted for Cd

(II), Cr (VI) and other cations by using green algae, brown algae and red algae

used in biosorption studies (Bishnoi et al., 2007; Lawal et al., 2007; Arief et al.,

2008; Gupta and Rastogi, 2008; Yahya et al., 2009; Pahlavanzadeh et al.,

2010; Rosales and Cruz, 2010). The data are significantly evident as can be

seen from the values of R2 in the present data. The biosorption capacity (qmax)

141

of the present study on Cd (II) ion biosorption by S. communis, C. delmatica,

and Spirogyra spp. do not show significant variations. These values are also

closer to those previously found in biosorption studies on Cd (II) ion by using

other green algae, such as Spirogyra insignis and Chlorella crispus, Ulva

lactuca (Romera et al., 2006; Aksu and Donmez, 2001; Sari and Tuzen, 2008).

However, these values show variation from those of the similar studies

conducted by using organisms belonging to the other groups, such as red

algae, brown algae, blue green algae, fungi, lichen, cotton, Hydrilla verticellata,

etc. (Romera et al., 2006; Liu et al., 2007; Doshi et al., 2007; Zafar et al., 2007;

Iqbal et al., 2007; Naddafi et al., 2007). Similarly, the biosorption capacity (qmax)

of the present study on Cr (VI) ion biosorption by S. communis, C. delmatica,

and Spirogyra spp. do not show significant variations. However, these values

show variation from those of the similar studies conducted by using sorbents,

such as Bacillus licheniformes, green taro (Colocasia esculenta),

Saccharomyces cerevisiae, tea factory wastes, water hyacinth, and water lilly

flower (Elangovan et al.,2008; Malkoc and Nuhoglu, 2007). Previously,

favourable biosorbents have been reported to have low Langmuir constant (b)

values and high ―qmax‖ values (Kratochvil and Volesky, 1998); the present study

confirms this assumption (Table 4.1). The similar biosorption capacity of Cd (II)

ion and Cr (VI) of various organisms belonging to the group of green algae

might be attributed to the biochemistry of cell wall (Davis et al., 2003c) and the

functional groups responsible which provide binding sites to cations. The FTIR

spectra (Figure 4.64a; Figure 4.64b; Figure 4.64c; Table 4.8), as its discussion

follows, provide an evidential support to this idea. ‖

5.1.6. Kinetic Studies of Biosorption

―Study of kinetic parameters is helpful in the prediction of the rate-

limiting step in the biosorption of the studied metal ions, Cd (II) and Cr (VI) onto

the algal surface under investigation, S. communis, C. delmatica, and

Spirogyra spp. For the solid-liquid biosorption, the solute transfer process can

be characterized by either external mass transfer or intra-particle diffusion or

both (Allen and Brown, 1995; Aravindhan et al., 2004). The biosorption

dynamics can be described by various successive steps, such as transport of

solute from bulk solution to the biosrbent‘s exterior surface, solute diffusion into

142

the pore of biosorbent, and biosorption of solute on the interior surfaces of the

pores of the biosorbent. The overall rate of biosorption will be controlled by the

slowest step, which would be either film diffusion or pore diffusion (Aravindhan

et al., 2004). The external mass transfer controls the biosorption process for

the systems that have poor mixing, dilute concentration of sorbate, small

particle sizes of biosorbent and higher affinity of sorbate for biosorbent.

Whereas, the intra-particle diffusion will control the biosorption process for a

system with good mixing, large particle sizes of biosorbents, high concentration

of sorbate, and low affinity of sorbate for biosorbent. ‖

―From the figures 4.2, 4.12, 4.22, 4.32, 4.44, 4.42, and 4.52, it was

observed that the rate of biosorption was fastest in the first 60 min., followed by

another 60 min. of average sorption rate and thereafter no significant metal

uptake was observed. This initial rapid phase of biosorption may be due to

higher number of vacant binding sites at the initial stage, as a result a high

concentration gradient between sorbate in solution and the sorbate on the

biosorbent surface may exist. As time proceeds, this concentration gradient

gets reduced due the accumulation of Cd (II) and Cr (VI) ions in the vacant

sites, leading to decreased biosorption rate at later stages. The principle behind

the biosorption kinetics involves the search for a best model that will represent

the experimental data. In the present investigations, the biosorption data were

analyzed using pseudo-second order model, as developed by Ho and Mckay

(1999), was applied for the interpretation of kinetic parameters of the

experimental data. This model is based on the assumption that the following

reaction proceeds as,‖

M+2 ions → biomass

―Where, ―M+2‖ is the metal ion (cation), biomass is the biosorbent, and

M+2 is being accumulated on the biosorbent and the rate of reaction can be

predicted by the following equation: ‖

qt = kqe2t /(1+kqet)

―Where ―k” is the pseudo second order rate constant (g/mg · min), qe is

the amount of metal ion adsorbed at equilibrium (mg/g), and ―qt‖ is amount of

cations [Cd (II) and Cr (VI)] on the algal surface (S. communis, C. delmatica,

and Spirogyra spp.) at any time (mg/g). This equation was linearized (equation

4.7) and plots of ‗t/q‘ versus ‗t‘ were obtained (Figures 4.8, 4.18, 4.28, 4.38,

143

4.48, 4.58 ). These plots clearly indicate that pseudo-second order model fits

very well to the available data. ‖

―Table 4.4, lists the rate constants studied for different initial metal ion

concentrations derived from the pseudo-second order models. The values of

K2, the second order rate constant (g/mg/min) are close to the experimental

values and that of regression coefficient, R2 was highly signicant (R>0.998)

confirming that the pseudo-second-order model best describes the kinetics of

the biosorption of Cd (II) and Cr (VI) ions. Low values of K2 shows the strong

affinity between metal ions [Cd (II) abd Cr (VI) ions] and the biosorbents, S.

communis, C. delmatica, and Spirogyra spp. (Fig 5.1). Therefore, the formation

of chemisorptive bonds (Fig. 5.2) involving sharing or exchange of electrons

between the sorbate and the sorbent may be considered as the rate-limiting

step of the process (Ho and McKay, 1999; Mashitah et al., 2008; Tsekova et

al., 2008). These results are in agreement with those from similar studies (Lodi

et al., 1998; Ho, 2006; Kumari et al., 2006; Gupta et al., 2008; Yahya et a.,

2009; lawal et al., 2010; Pahlavanzadeh et al., 2010; Tsekova et al., 2008). ‖

―Previous studies also reported that the equilibrium time of Cr(VI)

removal by seaweed and fungi varied from tens to hundreds of hours

depending on experimental conditions (Gupta et al., 2001; Salem and Allia,

2008; Dundar et al., 2008). Equilibrium time was also dependent on initial Cr

(VI) ion concentrations. The respective equilibrium time under initial Cr (VI)

concentrations of 50, 100, 150 and 250 mg L−1 was 30, 60 and 120 min.,

respectively (Figs. 4.12, 4.32, 4.52). This might be attributed to the gradual

appearance of Cr (III) in the solution with the removal of Cr (VI), indicating that

the Cr (VI) adsorbed on the algal biomass was reduced to Cr (III). The amounts

of Cr (VI) removed from the contaminated water were more than the amounts

of Cr (III) detected, suggesting that not all of the biosorbed Cr (VI) ions were

reduced to Cr (III), some of the reduced Cr (III) ions were released to the liquid

solution while some adsorbed on the biomass. The biosorption and bio-

reduction processes were likely to be physico-chemical transformation as the

biomass used in the present study was dried biomass and grounded into fine

particles, and might have lost its biological activity. At low pH, on the other

hand, Cr (VI) had a high redox potential and favored Cr (VI) bio-reduction

144

(Kratochvil et al., 1998; Han et al., 2006). In addition, reductants on the

biomass, such as carbohydrate and protein could supply electrons for Cr (VI)

bio-reduction, with partial release of soluble organics or ultimate oxidized

product, CO2 (Park et al., 2005). This explains why increasing biomass dosage

and lowering pH of the contaminated water could achieve more Cr (VI) removal

within a shorter period of time. ‖

5.1.7. Thermodynamic studies

―For engineering purpose, enthalpy (∆H), entropy (∆S), and Gibb‘s free

energy (∆G) factors generally need to be determined so that the spontaneity of

the process can be referred (Ho and Ofomaja et al., 2006; and Cho and Kim,

2003). The present study was conducted at different temperatures and

calculations were made to determine thermodynamic parameters, such as

(∆H), entropy (∆S), and Gibb‘s free energy (∆G) for the biosorption of Cd (II)

and Cr (VI) ions by S. communis, C. delmatica, and Spirogyra spp. (Tables, 4.5

and 4.6). In all the studies, values of ∆H and ∆S were found to be positive,

whereas values of ∆G to negative. This implies that adsorption of Cd (II) and Cr

(VI) ions were found to be spontaneous and feasible under the given

conditions. The elevated values of ∆S in biosorption of Cr (VI) by all the tested

species reflect the affinity of the biosorbents for Cr (VI) ions and suggest some

structural changes in biosorbent and biosorbate. The ∆G values also decrease

in magnitude on increasing the temperature up to 303 K in case of biosorption

of Cd (II) by S. communis, C. delmatica, and Spirogyra spp. and up to 313 K in

case Cr (VI) ion biosorption by all the tested species. Recently, a number of

studies concerning the effect of temperature on metal uptake and determination

of thermodynamics parameters, such as ∆H, ∆S and ∆G have been performed

by using various biosorbents (Dundar et al., 2008; Benguella and Benaissa,

2002; Malkoc and Nuhoglu, 2007; Gupta and Rastogi, 2008; Park et al., 2005;

Hassan et al., 2007; and Ho, 2006). These studies suggested that metal uptake

and feasibility of the reaction increases with increase in temperature (Aydin et

al., 2008; Yahya et al., 2009; Lawal et al., 2010). The present study also agrees

with these assumptions. Similar trends were also found in the present study by

using species of green algae, such as S. communis, C. delmatica, and

Spirogyra spp.‖

145

―The Arhenius constant (A₀) and Energy of activation (Ea) has also been

calculated in the present investigation. The values of high energy of activation

of Cd (II) biosorption by all the three tested biosorbents was found to be higher

than those of Cr (VI), indicating that Cd (II) ion biosorption involves

chemisorption processes. The Cd (II) sorption in the present study resembles

the previously performed investigation on Cu (II) uptake by fungus, Pynoporus

sanguiuneus, (Yahya et al., 2009) whereas Cr (VI) ion sorption showed

different values of A₀ and Ea. This might be attributed to the fact that Cr (VI) is

hexavalent, whereas Cd (II) and Cu (II) are divalent and the binding sites for

these ions and the mechanism of metal ion binding to the binding site may

vary. ‖

5.2. BINDING SITES FOR BIOSORPTION OF METAL IONS

―Fourier Transform Infra-red Spectroscopy (FTIR) and Scanning

Electron Microscopy (SEM) were carried out to determine the binding sites of

heavy metal ions, Cd (II) and Cr (VI). ‖

5.2.1. FTIR Spectroscopy

―Functional groups present on the surface of biosorbent are a very

important characteristic, which are largely characterized by the FTIR

spectroscopy method. This technique is only capable of providing a qualitative

description. Table 4.8 lists functional groups present in biosorbents under

investigation. As can be seen from figure 4.64, a shift in the peaks of the

spectra representing various functional groups was observed. Most of these

functional groups are electronegative, such as carboxyl, hydroxyl, carbonyl, etc

(Figure 4.66). Considering the biochemistry of call wall, It is quite predictable

that all these functional groups are components of organic constitutes, such as

aldehyde, ketone, carboxylic acid, alcohols, ethers, esters, etc., (Davis et al.,

2003c; Arief et al., 2008). ‖

―The present study further strengthens the already established concept

that these functional groups provide the binding sites responsible for

biosorption of cations on biosorbent surface. The most plausible explanation

146

behind the shift in terms of wavenumber for the –OH bending groups' contained

on the carboxylic acid was that these groups were included in the binding of

heavy metals, such as Cd (II) and Cr (VI) ion. C–O bond also seems to play a

vital role in cations binding. Similarly, stretching in the bonds of –C≒O and –C–

H also seem to be due to metal ion binding, such as Cr (VI), and Cd (II). ‖

―The present study has pointed out similar functional groups which have

been previously reported in other biosorbents (Davis et al., 2003a; Salem and

Allia, 2008; Dundar et al., 2008; Bueno et al., 2008; Xu and Liu, 2008; Gupta

and Rastgi, 2008; Yahya et al., 2009; and others).‖

―Specifically, biosorption studies with the help of FTIR have been helpful

to determine the availability of certain surface functional groups as part of the

structure of biosorbents and are capable of binding of toxic heavy metal ions,

such as Cd (II) and Cr (VI).‖

―Several previous studies have provided good knowledge of the metal

binding mechanism (Figure 4.67) on functional groups based on FTIR results

(Yu et al., 2000; Pavasant et al., 2006; Doshi et al., 2007; Elangovan et al.,

2008; Panda et al., 2008; Yu et al., 2000; Yazici et al., 2008). It is noticeable

that spectral shift in terms of wave number was mainly due to chemical metal

binding and the solution pH was not as responsible for the radical structural

change on the biosrbent surface as the chemical pretreatment. The solution pH

merely alters the -–OH bonds due to H+ ions binding to the functional groups.

Therefore, the identification of the functional groups in the present study is

extremely helpful to understand the surface binding mechanism of Cd (II) and

Cr (VI) ions from wastewater using bio-functional magnetic beads. ‖

5.2.2. Scanning Electron Microscopy (SEM)

―In the mechanistic study of Cd (II) ion and Cr (VI) removal by S.

communis, C. delmatica, and Spirogyra spp., SEM was utilized to acquire the

surface topology of unloaded and loaded biosorbent. Heavy metal-loaded

biomass differed in morphology compared to the unloaded samples. In the

case of loaded samples, the matrix layers of the cell wall were seen to shrink

and stick. This structural change was attributable to the strong cross-linking of

metal ions [Cd (II) and Cr (VI)] and negatively charged chemical groups on the

147

cell wall polymers. The mechanism of metal biosorption varies according to the

metal type and the type of biosorbent (Han et al., 2006). SEM is one of the

most reliable methods used to probe the surface complexation. The walls of

algal species under investigation had a plump shape with a transparent

external layer on the outer cell surface, which became wrinkled after the cell

adsorbed Cd (II) and Cr (IV) ions. This observation clearly indicated that the

primary Cr (VI) sequestering sites were on the cell wall surface instead of being

located on the intracellular sites. In case of Cd (II) ions adsorption, the sites

were found specifically located in a patch pattern, whereas in the case of Cr

(VI) biosorption, the metal ion seemed to bind onto the entire surface relatively

uniformly. The findings of the present study are in agreement with the

previously reported investigations by using other biosorbents, such as

Saragassum vulgaris, Chlorella miniata, yeast cells, etc (Raize et al., 2004;

yavuz et al., 2003; Das et al., 2007; etc.). Recently, Spirogyra spp., was used

for the removal of metal ion (lead) contaminate (Gupta and Rastogi, 2008); but

no comparison was shown between the loaded and unloaded biosorbents.

However, a number of studies previously conducted using SEM had provided

solid evidence to show biosorption of heavy metal contaminants (Zhou et al.,

2005; Kumari et al., 2006). ‖

5.2.3 Biochemistry of Biosorptive Surface

―The cell wall of green algae, such as S. communis, C. delmatica, and

Spirogyra spp. may consist of macromolecules, such as Cellulose in many (b-

1,4-glucopyroside), hydroxyproline glucosides; xylans and mannans etc.

(Davis et al., 2003a,c). Moreover, storage compounds, such as amylase and

amylopectin may also be present on the cell wall. These compounds possess

many functional groups, such as Carboxyl, hydroxyl, carbonyl, etc. These

functional groups are the biosorptive binding sites (already discussed in section

5.2.1) for the cations, such as Cd (II) and Cr (VI). ‖

5.2.4 Mechanism of Biosorption

―Various mechanisms have been proposed as mechanism of the studies

of biosorption by green algae, such as complexation, chelation, coordination,

148

ion exchange, precipitation, reduction, etc. However, in the present study some

of these seem to be more plausible as discussed below. ‖

―Ion exchange is a reversible chemical reaction where an ion within a

solution is replaced by a similarly charged ion attached onto an immobile solid

particle (Han et al., 2006). In general, the ion exchange mechanism can be

represented by the following equation:‖

―Here HY corresponds to the number of acid sites on the solid surface,

MX+ is metal ion, and MYX is the sorbed MX+. By considering the above

equation, the ion-exchange equilibrium constant can be determined by the

following equation (Arief et al., 2008). ‖

―On the basis of the above equation, ion exchange can be predicted as a

dominant mechanism for the cadmium ion biosorption onto S. communis, C.

delmatica, and Spirogyra spp. Previously, Romera et al., (2006) also claimed

that cadmium biosorption on de-alginated seaweed waste followed ion

exchange mechanism by using various techniques, such as potentiometric

titration, FTIR spectroscopy, calcium ion placement and esterification etc. Akar

et al. (2006) has also probed the interaction between cation and biomass and

declared biosorption process as expressed by ion exchange or complexation.

However, another plausible reason could be precipitation by phosphate

released from the biomass (also supported by the previous study, Cho and

Kim, 2003). The interaction between Cd (II) ions and the functional groups on

the surface of the algal cell wall may be also believed to occur by a

combination of ion exchange and complexation processes, as has been

proposed in the previous studies as well (Akar et al., 2006; Ofomaja and Ho,

2007). This phenomenon might involve the replacement of H+ ions on

functional groups of biomass by cadmium ions in the bulk solution (Asbchin et

al., 2008). The release of K+, Mg2+, Ca2+, and Na+ from the biosorbent during

heavy metal ion attachment may be considered as evidence to show the

involvement of ion exchange as dominant mechanism (Chen and Wang, 2006;

Fiol et al., 2006). Ion-exchange processes were also suggested as the main

biosorption mechanism by other studies (Xu and Liu, 2008; Panda et al., 2008;

149

Raize et al., 2004; Ofomaja and Ho, 2007; Dakiky et al., 2002). Considering the

biochemistry of the cell wall of the algal species under investigations, ionic

interactions, formation of complexes between metal cations and ligands

contained within the cell wall biopolymer structure and precipitation on the cell

wall matrix, may also be proposed. In previous studies of biosorption of Cd (II)

ion by the brown marine macroalga, Sargassum vulgaris, these ions have been

found to attach with chemical groups containing oxygen, carbon, nitrogen and

sulfur (Raize et al., 2004) showing that ion exchange might not be the main

mechanism for Cd (II) biosorption. Another frequently encountered metal

binding mechanism is chelation. Chelation can be properly defined as a firm

binding of a metal ion with an organic molecule (ligand) to form a ring structure.

As mentioned before, various functional groups including carboxylate, hydroxyl,

sulfate, phosphate amides, and amino groups can be considered responsible

for metal sorption. Among these groups, the amino group is the most effective

for removing heavy metals, since it does not merely chelate cationic metal ions

but also adsorbs anionic metal species through electrostatic interaction or

hydrogen bonding (Deng et al., 2007). ‖

―As far as Cr (VI) ion biosorption is concerned, at low values of pH ˂ 5.0,

the amino groups on the surface of algal cell wall are protonated and adsorb

anionic hexavalent chromium via electrostatic attractions, as shown by the

following equation.‖

―Where, (HVCr) M– represents anionic hexavalent chromium. In addition,

in the pH range of 2.4 to 7.0, both Cr (VI) and Cr (III) are present on the surface

of the biomass, suggesting that anionic hexavalent chromium ions were

reduced to Cr (III) during the sorption process. In turn, the resultant Cr (III)

species may then be chelated by the amine groups (Deng et al., 2007).

Furthermore, the importance of various functional groups on the surface of the

biomass, e. g., amino, phosphate, hydroxyl and carboxyl groups, may be

demonstrated by chemical modification. A reduction in adsorption capacity may

be observed upon modifying these chemical functional groups. The contribution

of these functional groups and the involvement of combination of chelation and

150

ion exchange mechanisms for Cr (VI) adsorption on biomass has also been

verified by the previous studies (Gode et al., 2008; Majumdar et al., 2008). ‖

――On the other hand, the phenomenon of physical adsorption cannot be

ruled out as far as biosorption of Cd (II) and Cr (VI) ions is concerned onto the

algal biomass under present investigation. In physical adsorption, a weak Van

der Waals attraction is observed between an adsorbate and a surface. From a

thermodynamic point of view, physical adsorption is spontaneous (negative

Gibbs energy) and exothermic (negative value of ∆H). Physical adsorption has

not yet been proved to be an important part in biosorption processes. However,

in several rare biosorption cases, physical adsorption may actually be predicted

as dominant. Recently, Cd (II) ion and other cations by using various

biosorbents, such as olive cake and rubber leaf powder (Al-Anber and Matouq,

2008; Ngah and Hanafiah, 2008) have also supported the phenomenon of

physical adsorption. ‖

Future directions:

―Based on the findings of the present investigations, a need arises for

the application of green algae at large scale to explore the various aspects

relevant to the industrial scale application of biosorption of toxic heavy metals.

The physicochemical conditions, for example, pH, multi-ionic composition may

be chosen to simulate the real wastewater on the basis of thermodynamics and

reaction kinetics studies. Optimization of the parameters of biosorption process

can also be done, including reuse and recycling by studying diffusion

resistance and fluid dynamics on biosorption column or chemical engineering

reactors, such as fluidized bed reactor etc. Improved mathematical modeling

may also be developed for better interpretations and prediction analysis in this

connection. Immobilization of biomaterials is another key aspect for the

purpose of biosorption application. It is important for decreasing the cost of

immobilization and consequently distribution, regeneration and reuse of

biosorbents (Tsezos, 2001; Hu and Wang, 2003). Application of hybrid

technologies may be introduced, which would involve combining various

processes to treat real effluent. Various biotechnology-based processes, such

as biosorption, bio-reduction and bio-precipitation may be suggested in this

context. ‖

151

―The above-mentioned bioprocesses along with other non-

biotechnology-based processes, such as chemical precipitation, flotation,

electrochemical processes, membrane technology, may also be helpful for

large scale wastewater treatment, for simultaneous removal of organic

substances and heavy metal ions from aqueous solution. A few studies like this

have been reported (Brady et al., 1994; Riordan et al., 1997; Thomas and

Macaskie, 1998). All these processes may possibly be realized in a single

reactor, hence the corresponding novel reactors should be designed (Tsezos,

2001). ‖

―To sum up, filamentous green algae, which have not been explored

earlier as has been done in this study, are a unique and promising biomass for

the biosorption of the toxic metals [Cd (II), Cr (VI), etc.] and is expected to

contribute to the bioremediation of polluted water and may also help to

elucidate the mechanism of removal of heavy metals using algal biomass.

Future study should be based on the advancements of biomass exploiting

processes following the well-tried and documented methodologies (Tsezos,

2001). ‖

152

Conclusions

―The capability of the use of three filamentous green algae, S.

communis, C. delmatica, and Spirogyra spp. were taken under investigations

for the removal of toxic metal ions, such as Cd (II) and Cr (VI). The findings

were based on biosorption capacity, equilibrium modeling, kinetic and

thermodynamic studies, and surface binding sites as observed by FTIR and

SEM. Experiments were performed as a function of initial solution pH,

temperature, initial metal ion concentration, biosorbent dosage and contact

time, etc. The solution pH, temperature and initial metal ion concentration

played a significant role in affecting the capacity of biosorption. Increase of the

pH over 5.0 (in case of divalent Cadmium) and 4.0 (in case of hexavalent

Chromium), temperature over 303 K and initial metal ion concentration of 2

mg/L in Cd (II) and 100 mg/L in Cr (VI) biosorption led to a reduction of the

biosorption capacity of the biomass. Optimum biosorbent dosage was 0.2 g/L

for Cd (II) ion and 2 mg/L for Cr (VI) ion in the metal ion solution. The

equilibrium concentrations of metal ion between the sorbate in the solution and

on the biosorbent surface were practically achieved in approximately 120 min

(depending on the temperature). Biosorption kinetics was found to follow

pseudo-second-order rate expression. Equilibrium biosorption data for all

biosorption studies were best represented by Langmuir and Temkin isotherms.

The typical dependence of metal ion uptake on temperature and kinetic studies

indicated the biosorption process of S. communis, C. delmatica, and Spirogyra

spp. for Cd (II) and Cr (VI) ions to be physical adsorption enhanced with

chemical effect and diffusion controlled. Thermodynamics parameters were

estimated for the biosorption of all the studies conducted. Negative values of

∆G indicate that biosorption is spontaneous and exothermic in nature. The

positive values of ∆S and ∆H reflect the affinity of the biosorbents for Cd (II)

and Cr (VI) ions. ‖

―The present study concludes that S. communis, C. delmatica, and

Spirogyra spp. may employed as a low-cost and eco-friendly biosorbents as an

alternative to the currently used expensive methods of removing toxic metal

153

ions, such as Cd (II) and Cr (VI) from industrial effluents. This may also help in

the development of subsequent novel technologies (with better efficiency and

eco-friendly) in the application of biosorption at industrial scale. ‖

154

CHAPTER 6

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