Republic of Iraq Ministry of Higher Education And Scientific Research Al-Nahrain University College of Science Department of Chemistry
ULTRAVIOLET VISIBLE INVESTIGATION OF SOME TRANSITION METAL COMPOUNDS IN
ROOM TEMPERATURE IONIC LIQUIDS MOLTEN SALTS
A thesis
Submitted to the College of Science
Al-Nahrain University
In partial fulfillment of Requirements
For the Degree of master in
Chemistry
BY Mariam Ali Saeed
B.SC. (AL-NAHRAIN UNIVERSITY 2005)
Supervised By Dr. Hadi M.A. Abood
May 2008 Rabeea Al-thani 1429
I
Supervisor�certification�
I certify that this thesis was prepared under my Supervision in the
Department of Chemistry, College of Science, and Al-Nahrain University as a
partial requirement for the degree of Master of Science in chemistry.
Signature:
Name: Dr. Hadi M.A. Abood
Date:
In view of the available recommendation, I forward this thesis for debate by
the Examining Committee.
Signature:
Name: Assistant Professor. Dr. Salman A. Ahmed
Head of Chemistry Department
College of Science
AL-Nahrain Univercity
II
Examining�Committee’s�Certification
We the examining committees, certify that we read this thesis and examined
the student Mariam Ali Saeed in its contents and that, according to our opinion,
is accepted as a thesis for the degree of master, in chemistry.
Signature:
Name:
(Chairman)
Signature: Signature:
Name: Name:
(Member) (Member)
Signature:
Name:
(Member\advisor)
Approved for the council of the College of Science, and for the College of
Graduate Studies.
Assistant Professor
Dr. Laith Abdul aziz Al-Ani
Dean of College of Science
Al-Nahrain University
III
Acknowledgement
First of all, thanks are due to God for his blesses. Then I would like to express my deepest gratitude and great appreciate to my respected supervisors, Dr.Hadi M.A. Abood for suggesting this project, and his valuable guidance, advice, support, and continuous encouragement through my study.
Sincere thanks to the chemistry department, Al-Nahrain University and to the staff of the department, specially the head of the department Dr. Salman A. Ahmed.
My deepest thanks and appreciations to our Professor Dr. Ayad H. Jassim (God bless his soul) for his help and support through all the years of my study.
My deepest thanks and appreciations to my lovely family, specially my mother & father (May God bless them) for their patient with me, and to all those who helped and support me through my work, I dedicate this work for them.
Finally, sincere thanks and deep respect goes to all my teachers, and friends for their help and support.
Mariam A. Saeed
2008
IV
List�of�Contents�
� Subject� � Pages�Thesis Title.
Supervisor Certification.
Examining Committee’s Certification.
Acknowledgement.
List of Contents.
List of Figures.
List of Tables.
List of Abbreviations.
Abstract.
Chapter One: - Introduction.
1.1. Introduction.
1.2. Room Temperature Ionic Liquids ‘RTILs’.
1.3. A short history of ionic liquids- from molten salt to
neoteric solvents.
1.4. Green chemistry and green solvents.
1.5. General characterization of ionic liquids.
1.6. Advantageous characterizations of RTILs.
1.7. Applications of Ionic liquids.
1.8. Room-Temperature Molten Salts Based on the
Quaternary Ammonium Ion.
1.9. General properties of Choline Chloride.
1.10. Deep eutectic solvents.
1.10.1. Eutectic mixtures of Urea and Choline chloride.
1.10.2. Eutectic mixtures of Malonic acid and Choline
chloride.
1.11. Interaction of the ligands with the metal ions.
1.12. The Chemistry of the first Transition Metal Elements.
I
II
III
IV
VIII
XIV
XVI
XVIII
1-42
1
2
2
5
6
8
12
13
14
15
15
19
21
25
V
1.12.1. Chromium (III), (d3).
1.12.2. Iron (III), (d5).
1.12.3. Cobalt (II), (d7).
1.12.4. Nickel (II), (d8).
1.12.5. Copper (II), (d9).
1.13. The coordination chemistry of transition metal cations
in chloride melts.
1.14. . Determination of spectra.
1.14.1. Orgel diagrams.
1.14.2. Tanabe-Sugano diagrams
1.15. Aim of this work.
Chapter Two: - Experimental Part (Materials and Methods).
2.1. Chemicals and Instruments.
2.1.1. Chemicals.
2.1.2. Instruments.
2.1.2.1. Electronic absorption spectra.
2.2. Setting of heating apparatus.
2.3. Drying method.
2.4. Preparation of room temperature ionic liquid.
2.4.1. Preparation of Choline chloride/Urea room
temperature ionic liquid.
2.4.2. Preparation of Choline chloride/Malonic acid room
temperature ionic liquid.
2.5. Preparation of metal chloride ionic liquid mixtures.
2.6. Equations.
2.7. Example of calculations.
Chapter Three: - Results and Discussion.
3.1. Introduction.
3.2. Metal chloride in room temperature ionic liquid mixtures.
25
27
29
34
36
38
39
39
40
42
43-52
43
43
44
44
44
44
45
45
45
45
46
48
52-106
52
53
VI
3.2.1. Spectroscopic investigation of Chromium (III)
chloride solution.
3.2.1.1. Chromium (III) chloride in choline
chloride/urea room temperature ionic
liquids.
3.2.1.2. Substitution Reaction with Ammonium
Nitrate.
3.2.1.3. Chromium (III) chloride in Choline
chloride/Malonic acid room temperature
ionic liquid solution.
3.2.2. Spectroscopic investigation of Iron (III) chloride
solution.
3.2.2.1. Iron (III) chloride in Choline
chloride/Urea room temperature ionic
liquid solution.
3.2.2.2. Substitution Reaction with Ammonium
Nitrate.
3.2.2.3. Iron (III) chloride in Choline
chloride/Malonic acid room temperature
ionic liquid solution.
3.2.3. Spectroscopic investigation of Cobalt (II) chloride
solution.
3.2.3.1. Cobalt (II) chloride in Choline
chloride/Urea room temperature ionic
liquid solution.
3.2.3.2. Substitution Reaction with Ammonium
Nitrate.
3.2.3.3. Cobalt (II) chloride in Choline
chloride/Malonic acid room temperature
53
53
57
59
64
64
66
67
72
72
77
80
VII
ionic liquid solution.
3.2.4. Spectroscopic investigation of Nickel (II) chloride
solution.
3.2.4.1. Nickel (II) chloride in Choline
chloride/Urea room temperature ionic
liquid solution.
3.2.4.2. Substitution Reaction with Ammonium
Nitrate.
3.2.4.3. Nickel (II) chloride in Choline
chloride/Malonic acid room temperature
ionic liquid solution.
3.2.5. Spectroscopic investigation of Copper (II) chloride
solution.
3.2.5.1. Copper (II) chloride in Choline
chloride/Urea room temperature ionic
liquid solution.
3.2.5.2. Substitution Reaction with Ammonium
Nitrate.
3.2.5.3. Copper (II) chloride in Choline
chloride/Malonic acid room temperature
ionic liquid solution.
3.3. Conclusion.
3.4. Suggestion for future work.
References.
83
83
87
90
96
96
98
100
104
106
107
�
�
�
VIII
List�of�Figures�
Figures� Pages�Chapter One: - Introduction. Figure (1-1):- Examples of simple room-temperature ionic
liquids.
Figure (1-2):- Examples of room temperature molten salts.
Figure (1-3):- Air and water stable molten salts.
Figure (1-4):- Some types of room temperature ionic liquid.
Figure (1-5):- Freezing point of Choline chloride/Urea
Mixture as a function of composition.
Figure (1-6):- Conductivity and viscosity of a 2 urea to 1
choline chloride mole ratio ionic liquid as a
function of temperature.
Figure (1-7):- Freezing points of Choline chloride with
phenylpropionic and phenylacetic acids as a
function of composition.
Figure (1-8):- Freezing points of Choline chloride with
Oxalic, Malonic, and Succinic acids as a function of
composition.
Figure (1-9):- Plot of log viscosity vs. reciprocal of
temperature for the Choline chloride/Malonic acid
ionic liquid.
Figure (1-10):- Spectra of CrCl3 acetamide-potassium nitrate
at 90oC and in butyramide-sodium nitrate at 140oC.
Figure (1-11):- Spectra of Co (II) in LiCl/KCl eutectic.
Figure (1-12):- Absorption spectra of Co (II) in AICI3,-BPC
melt at 40.0 OC.
Figure (1-13):- Spectra of CoCl2 in acetamide-KI at different
6
7
8
13
17
18
20
20
21
27
32
32
33
IX
temperature (90 and 140oC) and of butyramide-
sodium nitrate at 140oC.
Figure (1-14):- Spectra of NiCl2, in acetamide-potassium
nitrate at 90oC and in butyramide-sodium nitrate at
140oC.
Figure (1-15):- Copper (II) absorption spectra in solvents
containing chloride ion.
Fig.ure (1-16):- The splitting of free-ion D terms in
octahedral and tetrahedral fields.
Figure (1-17):- The splitting of free-ion F and P term in
octahedral and tetrahedral field.
Chapter Two: -Experimental Part. Figure (2-1):- Splitting of terms for dn configurations in to (a)
and (b).
Chapter Three: -Results and Discussion. Figure (3-1):- The U.V.-Vis spectrum of Chromium (III) of
(7.4×10-2m) in Choline chloride/Urea room
temperature ionic liquid.
Figure (3-2):- Crystal field splitting of 4F term of d3 ion
octahedral configuration.
Fig. (3-3):- Tanabe-Sugano diagram for d3 ion configuration.
Fig. (3-4):- The U.V.-Vis. Spectra of CrCl3.6H2O (0.1 M) in
aqueous solution (curve A), and in choline
chloride/ urea eutectic solution (7.4×10-2m), (curve
B) at room temperature.
Figure (3-5):- The U.V.-Vis. Spectra of dissolved CrCl3.6H2O
ion in choline chloride/urea of (7.4×10-2m) (curve
A), and in choline chloride/urea with addition of
NH4NO3 salt (curve B) eutectics at room
35
38
40
40
46
54
54
55
56
58
X
temperature.
Figure (3-6):- The U.V.-Vis. Spectrum of CrCl3.6H2O ion of
(7.5×10-2m) in Choline chloride/ Malonic acid
room temperature ionic liquid.
Figure (3-7):- The U.V.-Vis. Spectra of CrCl3.6H2O ion in
choline chloride/urea of (7.4×10-2m) (curve A), and
in choline chloride/malonic acid of (7.5×10-2m)
(curve B) eutectics at room temperature.
Figure (3-8):- Suggested structures for Chromium (III) in (1)
choline chloride/urea ionic liquid, (2) choline
chloride/urea ionic liquid with added salt NH4NO3
and (3) choline chloride/ malonic acid ionic liquid
all formed at room temperature.
Figure (3-9):- The U.V.-Vis. spectrum of pale yellow solution
of FeCl3.4H2O of (6.6×10-2m) in Choline
chloride/Urea ionic liquid at room temperature
after 24 hours.
Figure (3-10):- The U.V.-Vis. spectrum of brownish yellow
solution of FeCl3.4H2O of (6.6×10-2m) in Choline
chloride/Urea ionic liquid at room temperature
after 72 hours.
Figure (3-11):- The U.V.-Vis. spectrum of FeCl3.4H2O of
(6.6×10-2m) in Choline chloride/Urea with addition
of NH4NO3 salt room temperature ionic liquid.
Figure (3-12):- The U.V.-Vis. spectrum of FeCl3.4H2O of
(6.9×10-2m) Choline chloride/Malonic acid ionic
liquid at room temperature.
Figure (3-13):- Tanabe-Sugano diagram for Octahedral and
Tetrahedral for d5 ion configuration.
60
61
63
65
65
66
68
86
XI
Figure (3-14):- Suggested structures of Iron (III) complexes
in (1) choline chloride/ urea ionic liquid, (2)
choline chloride/urea ionic liquid with added
NH4NO3 salt, (3) choline chloride/malonic acid
ionic liquids all formed at room temperature.
Figure (3-15):- The U.V.-Vis spectrum of CoCl2.6H2O of
(6.9×10-2m) Choline chloride/Urea room
temperature ionic liquid.
Figure (3-16):- The U.V.-Vis. spectrum of the diluted solution
of CoCl2.6H2O of (3.45×10-2m) in Choline
chloride/Urea room temperature ionic liquid.
Figure (3-17):- Crystal field splitting of 4F term of d7 ion
Tetrahedral configuration.
Figure (3-18):- Tanabe-Sugano diagram for d7 Tetrahedral
ionic configuration.
Figure (3-19):- The U.V.-Vis. spectra of Cobalt (II) in
Choline chloride/Urea room temperature ionic
liquid after 40 days, (curve A) concentrated
solution (6.9×10-2m) and (curve B) diluted solution
(3.45×10-2m).
Figure (3-20):- The U.V.-Vis. spectrum of CoCl2.6H2O of
(6.9×10-2m) in Choline chloride/Urea room
temperature with addition of NH4NO3.
Figure (3-21):- The U.V.-Vis. spectra of CoCl2.6H2O of
(3.45×10-2m) in Choline chloride/Urea room
temperature ionic liquid (curve A) and with
addition of NH4NO3 (curve B).
Figure (3-22):- The U.V.-Vis. spectrum of CoCl2.6H2O of
(7.4×10-2m) in Choline chloride/Malonic acid room
71
72
73
74
75
76
78
79
81
XII
temperature ionic liquid.
Figure (3-23):- Suggested structures for Cobalt (II) ions in (1)
choline chloride/urea ionic liquid, (2) choline
chloride/urea ionic liquid with added salt NH4NO3
and (3) choline chloride/malonic acid ionic liquid
all formed at room temperature.
Figure (3-24):- The U.V.-Vis. spectrum of NiCl2.6H2O of
(1.1×10-1m) in choline chloride/ urea room
temperature ionic liquid.
Figure (3-25):- Crystal field splitting of 3F term of d8 ion
Octahedral configuration.
Figure (3-26):- Tanabe-Sugano diagram for d8 Octahedral
electronic configuration.
Figure (3-27):- The U.V.-Vis. spectra of NiCl2.6H2O in
aqueous solution of 0.1 M (curve A) and in choline
chloride/ urea ionic liquid of (1.1×10-1m) at room
temperature.
Figure (3-28):- The U.V.-Vis. spectrum of NiCl2.6H2O in
choline chloride/ urea with addition of NH4NO3
ionic liquid at room temperature.
Figure (3-29):- The U.V.-Vis. spectra of NiCl2.6H2O in
choline chloride/urea with addition of NH4NO3
ionic liquid (curve A) and pure melt of choline
chloride/ urea (curve B) at room temperature.
Figure (3-30):- The U.V.-Vis. Spectrum of NiCl2.6H2O of
(1.2×10-1m) in choline chloride/ malonic acid ionic
liquid at room temperature.
Fig. (3-31):- Tanabe-Sugano diagram for d8 tetrahedral
electronic configuration.
82
84
84
86
87
88
89
91
92
XIII
Figure (3-32):- Suggested structures for Nickel (II) ions in (1)
choline chloride/ urea ionic liquid, (2) choline
chloride/urea ionic liquid with added salt of
NH4NO3 and (3) choline chloride/ malonic acid
ionic liquid all formed at room temperature.
Figure (3-33):- Crystal field splitting for 2D configuration.
Figure (3-34):- The U.V.-Vis. Spectrum of CuCl2.2H2O of
(7.7×10-2m) in choline chloride/urea ionic liquid at
room temperature.
Figure (3-35):- The U.V.-Vis. spectrum of CuCl2.2H2O in
choline chloride/ urea ionic liquid with addition of
NH4NO3 at room temperature.
Figure (3-36):- The U.V.-Vis. spectra of CuCl2.6H2O in
choline chloride/ urea with addition of NH4NO3
ionic liquid at room temperature (curve A), and in
pure melt (curve B).
Figure (3-37):- The U.V.-Vis. spectrum of CuCl2.6H2O of
(8.03×10-2m) in choline chloride/ malonic acid
ionic liquid at room temperature.
Figure (3-38):- Suggested structures for Copper (II) ions in
(1) choline chloride/ urea ionic liquid, (2) choline
chloride/urea ionic liquid with added salt of
NH4NO3 and (3) choline chloride/malonic acid
ionic liquid all formed at room temperature.
95
96
97
99
100
101
103
�
�
�
XIV
List�of�Tables�
Tables� Pages�Table (2-1):- The chemicals used in this work with their
purities and suppliers.
Table (2-2):- Concentration of metal chloride in aqueous
solution and room temperature ionic liquid of
choline chloride/urea and choline chloride/malonic
acid.
Table (2-3):- Splitting of terms for dn configuration.
Table (2-4):- Values of (g) and (f) factors for metal ions and
various ligands.
Table (2-5):- Peak position of Chromium (III) ions in room
temperature ionic liquid of Choline chloride/Urea.
Table (3-1):- Electronic spectra for complexes of Chromium
(III) ions in aqueous solution and various ionic
liquids.
Table (3-2):- The molar absorptivity of Chromium (III) ions in
ionic liquid based on choline chloride.
Table (3-3):- Electronic spectra of complexes of Iron (III)
ions in various ionic liquids.
Table (3-4):- The molar absorptivity of Iron (III) ions in ionic
liquid based on choline chloride.
Table (3-5):- Electronic spectra for complexes of Cobalt (II)
chloride in various room temperature ionic liquids.
Table (3-6):- Electronic spectra of complexes of Cobalt (II)
ions in aqueous and various ionic liquids.
Table (3-7):- The molar absorptivity of Cobalt (II) ions in
ionic liquid based on choline chloride.
43
46
47
47
48
62
64
70
71
77
82
83
XV
Table (3-8):- Electronic spectra of complexes of Nickel (II)
ions in aqueous solution and ionic liquids.
Table (3-9):- Electronic spectra of complexes of Nickel (II)
ion in aqueous solution and various ionic liquids
solution.
Table (3-10):- The molar absorptivity of Nickel (II) ions in
ionic liquid based on choline chloride.
Table (3-11):- The ultraviolet visible spectroscopic behavior
of Copper (II) in various solvents.
Table (3-12):- Electronic spectra of complexes of Copper (II)
ions in choline based ionic liquids at room
temperature.
Table (3-13):- The molar absorptivity of Copper (II) ions in
ionic liquid based on choline chloride.
85
94
95
98
102
102
XVI
List�of�Abbreviations�
Abbreviation Name
UV-Vis Ultraviolet-Visible
Oh Octahedral
Td Tetrahedral
B Racah parameter
β Nephelauxetic factor
nm Nanometre
RTILs Room temperature ionic liquids
TFA Trifluoroacetate
Tf2N Bis(trifluoromethylsulfonyl) imide
PF6 hexafluorophosphate
EEIm 1,3-Diethylimidazolium
BMIm 1-Butyl-3-methylimidazolium
EMIm 1-Ethyl-3-methylimidazolium
BF4 tetrafluoroborate
[BMMIm]Tf2N 1-Butyl-2,3-dimethylimidazolium
bis(trifluoromethylsulfonyl) imide
Acac Acetylacetonate
phen 1,10-Phenanthroline
en Ethylenediamine
bipy 2,2-Bipyridine
PPh3 Triphenylphosphine
Dppe 1,2-Bis(diphenylphosphino)
ethane
VOC Volatile Organic Compounds
XVII
DMSO Dimethylsalfoxide
BEIm 1-butyl,3-ethylimidazolium
(CF3SO2)2N- Bis(trifluoromethyl)sulphomide
anion
M Concentration in unit of (mol/l)
‘Molarity’
m Concentration in unit of (mol/Kg)
‘molality’
BPC 1-n –butylpyridinium chloride
MEIC 1-methyl-3-ethylimidazolium chloride
[BMP]Tf2N 1-Butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl) imide
[(CH3)3N+CH2CH2OH.2(NH2)2C2O42.-
2(NH2)2CS (2-hydroxy-ethyl)trimethyl ammonium oxalate dithiourea
HOCH2CH2N+(CH3)3F-.(2NH2)2 (2-hydroxy-ethyl)trimethyl ammonium fluoride diurea
CFT Crystal field theory MOT Molecular orbital theory
XVIII
Abstract
Investigation of some transition metal complexes were carried out at room
temperature in the new clan of room temperature ionic liquids of choline
chloride based composition.
The Ultraviolet Visible Spectroscopic spectra of Cr3+, Fe3+, Co2+, Ni2+ and Cu2+
indicated a complexes formation between the metal ions and the entities of these
types of RTILs solvents namely choline chloride/ urea and choline chloride/
malonic acid which have freezing point at 12oC and 0oC respectively.
The complexes in choline chloride/ urea were found to be of Octahedral
geometry for Cr3+, Ni2+ and Cu2+, and Tetrahedral geometry for Fe3+, Co2+ metal
ions. Their spectra were found to be assigned to normal transition state of
corresponding dn configuration in both geometries of Octahedral and
Tetrahedral.
Substitution of solvent coordinated to these metals by added nitrate ions were
observed, suggested a weak bonding between solvent molecules and transition
metal cations, although they showed some geometrical coordination of
corresponding metal cations. All complexes were found to reach final state of
coordination at 72 hours, except Co2+ complexes when they showed instability
when they were left at room temperature for 40 days.
In choline chloride/ malonic acid the geometrical coordination of Ni2+ and
Cu2+ showed tetrahedral geometry rather than octahedral geometry in this RTIL,
while the other cations have similar coordination geometry.
CFT splitting value (10Dq) have been determined for most of new formed
complexes using appropriate dn Orgel and Tanabe-Sugano diagrams, and
Jorgenson rule was used to asset the suggestion of the type of ligands
coordinated to transition metal cations. The nature of the bonding between the
XIX
metal ions and the donor ligands were demonstrated through the Jorgenson rule
and calculation of Racah parameters B were obtained form appropriate dn
Tanabe-Sugano diagrams together with the qualitative assistance of Orgel
diagrams.
CHAPTER ONE
INTRODUCTION
CHAPTER TWO
EXPAREMANTAL
Part
(MATERALS &
METHODS)
CHAPTER THREE
RESULTS &
DISCUSSION
REFFRENCES
Chapter One Introduction
1
Chapter One
Introduction
1.1. Introduction.
What is Molten Salt? Molten salts are evidently numerous and potentially most significant and useful
class of nonaqueous solvents. While “salts” is widely used though imprecise term,
this class can be broadly defined as the liquid state of those compounds which
melt, usually above room temperature to give liquids displaying ionic properties, at
least to some degree. While they are probably more numerous than other
nonaqueous solvents, molten salts are currently the least-known and least
appreciated of the nonaqueous solvents, certainly much less so than their practical
importance demands.(1,3)
Molten Salt is a rather dreadful name for an otherwise useful category of
materials and processes. The term "Molten Salt" is self-descriptive; it is melted
salt(s). Another common name is Fused Salt(s).(2,4)
Salts are simple, usually ionic (that is the chemical bonds are a simple ionic
type) and stable compounds. The most common example of molten salt is "table
salt", or sodium chloride (NaCl). Both sodium and chlorine are reactive; sodium is
one of the most electropositive substances (tends to lose an electron) and chlorine
one of the most electronegative (tends to take an electron). These two opposite
substances readily join to form stable sodium chloride via a strong ionic bond. The
melting point of sodium chloride is 801° C. At which point it becomes a liquid, and
thus a "molten salt". This liquid is stable, has a heat capacity similar to water (by
volume) and flows much like water does. The major differences are the obvious
higher temperatures attainable in the molten salt state and when the salt solidifies
(freezes) it contracts versus expanding in water. Thus, molten salt freezing in a
pipe would not burst the pipe as water would.(3)
Chapter One Introduction
2
Molten salts are of considerable interest, not only from the technological point
of view, but also from the standpoint of satisfying a fundamental curiosity.
Structural studies of molten salts (such as alkali halides) by x-ray and neutron
diffraction have been undertaken as part of a general program to determine the
properties of molten salts.(3)
Molten salts, or if they are liquid under ordinary conditions, ionic liquids, are
interesting for the special properties unique to fluids of charged particles and for
applications of ionic liquids as environmentally friendly solvents.(3)
1.2. Room Temperature Ionic Liquids ‘RTILs’. Ionic liquids are highly solvating, non-coordinating medium in which a variety
of organic and inorganic solutes are able to dissolve. They are outstanding good
solvents for a variety of compounds, and their lack of a measurable vapour
pressure makes them a desirable substitute for VOCs. Ionic liquids are attractive
solvents as they are non-volatile, non-flammable, have a high thermal stability and
are relatively inexpensive to manufacture. They usually exist as liquids well below
room temperature up to a temperature as high as 200oC.(5,6)
An ionic liquid is a liquid that contains essentially only ions. Some ionic liquids,
such as ethyl ammonium nitrate, are in a dynamic equilibrium where at any time
more than 99.99% of the liquid is made up of ionic rather than molecular species.
In the broad sense, the term includes all molten salts, for instance, sodium chloride
at temperatures higher than 800 °C. Today, however, the term "ionic liquid" is
commonly used for salts whose melting point is relatively low (below 100 °C). In
particular, the salts that are liquid at room temperature are called room-temperature
ionic liquids, or RTILs.(7,8)
1.3. A short history of ionic liquids—from molten salts to neoteric
solvents.
Ionic liquid is a salt with a melting temperature below the boiling point of
water. Most salts identified in the literature as ionic liquids are liquid at room
temperature, and often to substantially lower temperatures.The history of ionic
Chapter One Introduction
3
liquids may be viewed as a relatively recent one, or one extending back to the 19th
century. Which history you choose depends on your definition of ‘ionic liquid’ and
how deeply you look for progenitors of the present materials receiving so much
interest for green chemistry applications.(9) Accordingly, this may be the birth of
new and important industrial materials, or we are witnesses to a renaissance of
molten salt chemistry. One fairly general feature of ionic liquids that is not part of
the definition is that most have organic cations and inorganic anions. There are
many synonyms used for ionic liquids that complicate a literature search.(10)
‘Molten salts’ is the most common and most broadly applied term for ionic
compounds in the liquid state. Unfortunately the term ‘ionic liquids’ was also used
to mean the same thing long before there was much literature on low melting salts.
It may seem that the difference between ionic liquids and molten salts is just a
matter of degree (literally), however the practical differences are sufficient to
justify a separately identified niche for the salts that are liquid around room
temperature.(11) Synonyms for materials that meet the working definition of ionic
liquid are ‘room temperature molten salt,’ ‘low temperature molten salt’, ‘ambient
temperature molten salt’, ‘ionic fluid’ and ‘liquid organic salt.’ The term ‘neoteric
solvent’ has been used in recent years to mean new types of solvents, or older
materials that are finding new applications as solvents. Supercritical fluids and
ionic liquids are good examples of neoteric solvents that have been known for a
long time, but are being considered as process solvents.(10)
The link between ionic liquids and green chemistry is clearly related to the
solvent properties of ionic liquids. Liquids range is the span of temperatures
between the freezing point and boiling point of a liquid. The consequence for green
chemistry is that ionic liquids are the ultimate non-volatile organic solvent, with
emphasis on the ‘non-‘.(12) No molecular solvent (other than molten polymers)
comes even close to the low volatility of ionic liquids. Ionic liquids have many of
the advantages of molten salts, and they avoid the worst disadvantages caused by
Chapter One Introduction
4
high temperature. In fact they are better viewed as a new type of non-aqueous
solvent, which is a neoteric solvent.(13,14)
The early history of ionic liquids began in 1914 when the first report of a room
temperature molten salt was reported by Walden. He reported the physical
properties of ethyl ammonium nitrate, [C2H5NH3] NO3, which has a melting point
of 12oC, formed by the reaction of ethylamine with concentrated nitric acid. Then,
Hurley and Weir stated that a room temperature ionic liquid could be prepared by
mixing and warming 1-ethylpyridinium chloride with aluminum chloride(14). In
1970s and 1980s, Osteryoung et al(19). and Hussey et al.(20) respectively carried out
extensive research on organic chloride–aluminum chloride ambient temperature
ionic liquids and the first major review of room temperature ionic liquids was
written by Hussey.(20) It was about in the middle of the 1990s when it was decided
in the community to replace the term ‘‘room temperature molten salt’’ by ‘‘ionic
liquid’’, and an ionic liquid is defined today as a liquid consisting solely of cations
and anions with a melting point of 100oC and below. Although any high
temperature molten salt is an ionic liquid, too, this novel term for the room
temperature liquids clearly made a distinction.(15,16)
The ionic liquids based on AlCl3 can be regarded as the first generation of ionic
liquids. The hygroscopic nature of AlCl3 based ionic liquids has delayed the
progress in their use in many applications since they must be prepared and handled
under inert gas atmosphere. Thus, the synthesis of air and water stable ionic
liquids, which are considered as the second generation of ionic liquids, attracted
further interest in the use of ionic liquids in various fields. (17) Generally, these
ionic liquids are water insensitive; however, the exposure to moisture for a long
time can cause some changes in their physical and chemical properties.(19) In 1992,
Wilkes reported the first air and moisture stable ionic liquids based on 1-ethyl-3-
methylimidazolium cation with either tetrafluoroborate or hexafluorophosphate as
anions. Unlike the chloroaluminate ionic liquids, these ionic liquids could be
prepared and safely stored outside of an inert atmosphere (18).
Chapter One Introduction
5
1.4. Green chemistry and green solvents. Green chemistry involves the development of chemical products and synthetic
procedures, which are environmentally friendly and have reduced health risks with
the search for more efficient methods to do chemistry. Its roots stem back some
years from a simple idea to a prominent concept which permeates all areas of
modern chemistry.(20) Green solvents are environmentally friendly solvents or
biosolvents, which are derived from the processing of agricultural crops. The uses
of petrochemical solvents are the key to the majority of chemical processes but not
without severe implications on the environment. The Montreal Protocol identified
the need to re-evaluate chemical processes with regard to their use of volatile
organic compounds (VOCs) and the impact theses VOCs has on the environment.
Green solvents were developed as a more environmentally friendly alternative to
petrochemical solvents, Ethyl lactate, for example is a green solvent derived from
processing corn. Lactate esters solvents are commonly used solvents in the paints
and coatings industry and have numerous attractive advantages including being
100% biodegradable, easy to recycle, non-corrosive, non-carcinogenic and non-
ozone depleting.(21,22)
Room temperature ionic liquids which melt below 100oC can be achieved by
for example incorporating a bulky asymmetric cation into the structure to stop the
ions packing easily. Since the melting points are low, ionic liquids can act as
solvents in which reaction can be preformed, and because the liquid is made of
ions rather than molecules, such reactions often give distinct selectivity and
reactivity when compared with conventional organic solvents. Room temperature
ionic have been utilized as clean solvents and catalysts for green chemistry and as
electrolytes for batteries, photochemistry and electro-synthesis.(23)
These stemmed significantly as ionic liquids have no significant vapor pressure
and thus create no volatile organic components. They also allow for easy
separation of organic molecules by direct distillation without loss of the ionic
liquid. Their liquid range can be as large as 300o C allowing for large reaction
Chapte
kinetic
reactor
[CF4]-, [
remarka
1.5. G Ionic
has a lo
the crys
come in
cation)
example
chloride
contain
depend
dialkylim
simple
melting
compos
archetyp
methylim
er One
control, w
volumes t
[PF3]-, [CF
ably high t
General liquids co
ow degree
stalline fo
n two mai
and bina
e, [EtNH3
e and 1,3
several d
upon th
midazoliu
salt are g
behavior
ition, and
pal sy
midazoliu
Fig.
which, cou
to be utili
F3CO2]-,[C
thermal st
characteonsist of a
e of symm
rm of the
in categor
ary ionic
3][NO3] i
3-dialkylim
different
he mole
um chlorid
given in F
r. For th
d this com
ystem,
um).(25,26)
.(1-1):- Exam
upled wit
ized. Salt
CF3SO3]-,
tability.(22,
erizationa salt wher
metry. The
e salt, and
ries, name
liquids (s
is a simp
midazolium
ionic spe
fractions
de presen
Figure (1-
he binary
mplex beh
[emim]
mples of sim
6
h their go
based up
etc, are w24)
n of ionicre one or b
ese factors
d hence lo
ely simple
salts whe
ple salt w
m chlorid
cies, and
s of the
nt. Examp
-1) and F
systems,
havior ha
]Cl-AlCl3
mple room-te
ood solve
pon poor n
water and
c liquids.both the io
s tend to r
ower the m
e salts (m
ere equilib
whereas m
des (a bi
their me
aluminu
ples of ion
Figure (1-
, the me
as been s
([em
emperature
nt proper
nucleophil
air insens
ons are lar
reduce the
melting po
made of a
brium is
mixtures
inary ioni
elting poin
um(III) ch
nic liquid
-2) and th
lting poin
studied ex
mim]+
ionic liquid
Introd
rties, allow
lic anions
sitive and
rge, and th
e lattice en
oint. Ionic
single an
involved)
of alumin
ic liquid
nt and pr
hloride a
ds consisti
hese show
nt depend
xtensively
= 1-
ds.(25)
duction
ws small
such as
d possess
he cation
nergy of
c liquids
nion and
).(24) For
num(III)
system)
roperties
and 1,3-
ing of a
w simple
ds upon
for the
-ethyl-3-
Chapte
With
solvents
polar or
extractio
employe
health a
impetus
the 1,3-
vapour p
as a new
synthesi
use of i
chemica
many re
processe
solubilit
it leaves
Room
alternati
liquids
crystalli
chloride
mp=80o
chloride
er One
F
h many c
s in chemi
rganic solv
on techno
ed historic
and enviro
s to search
-dialkylim
pressure, l
w class of
is and cata
ionic liqui
al process,
eactions ca
es taking
ty in the w
s the proce
m tempera
ive, ‘green
can also
ine structu
e, mp=80o C) or
e/potassium
Fig.(1-2):- E
chemical t
istry is vit
vents have
ologies an
cally. How
onmental e
h for ‘gree
midazolium
low viscos
f ‘green’ s
alysis. Th
ids in gre
, from an
an be perf
place in
water phas
ess, both f
ature molt
n’ technol
be descr
ures. This
01oC), or
even eut
m chlorid
Examples of
transform
al. Althou
e been ext
nd have l
wever, the
effects of
en’ techno
m cation, w
sity and h
solvents fo
he main m
een techno
ecologica
formed wit
solution. N
se, and the
for econom
ten salts (
logy in bo
ibed as l
definition
rganic co
tectic mix
e, 6/4, mp
7
room tempe
mations tak
ugh, any li
tensively u
largely su
e ever-incr
some org
ologies Ion
with a lar
igh therm
or extracti
milestones
ologies.(27)
al point of
thout solv
Note that
erefore mu
mic and en
(ionic liqu
th organic
iquid com
n involves
ompounds
xtures of
p=352o C)
erature molt
king plac
iquid coul
used for b
uperseded
reasing aw
ganic solve
nic liquids
rge range
mal and che
ion and se
reached in) The obv
f view, inv
vent, chem
most of
ust be stri
nvironmen
uids) are
c and orga
mpounds t
s pure ino
s (tetrabu
f inorgani
) (28) or or
ten salts.(25,2
e in solu
d be empl
oth synthe
water, w
wareness
ents has g
s, especial
of liquid
emical stab
eparation p
n the last
viously ide
volves no
mistry rema
solvents h
ipped from
ntal reason
emerging
anometalli
that displ
organic com
utylphosph
ic salts (
rgano-min
Introd
26)
ution, the
loyed as a
etic chemi
which was
of the det
given chem
lly those b
phase, ne
bility are
processes,
two years
eal situati
solvent. A
ains domin
have a sig
m the wate
ns.(26,29)
g as an at
ic chemist
lay ionic-
mpounds
honium c
(such as
neral comb
duction
role of
a solvent,
istry and
s widely
trimental
mists the
based on
egligible
emerged
, organic
s, on the
on for a
Although
nated by
gnificant
er before
ttractive,
try. Ionic
covalent
(sodium
chloride,
lithium
binations
Chapte
(triethyl
chloroal
N-alkylp
chloride
These
for high
synthesi
particula
phospha
to a broa
Ionic
have rec
These u
books,
liquids.
and high
1.6. A RTI
suitable
non app
1. Lo
coordina
er One
lammoniu
luminate i
pyridinium
e (Figure 1
e liquids h
hly charge
is and for
ar 1-n-b
ate (BMI.B
ader spect
liquids a
ceived gre
unique pro
these are
The mos
h ion cond
AdvantagILs posse
in numer
plicable or
w nucleop
ating envi
um chlor
ionic liqu
m chlorid
1-1 and 1-
have been
ed comple
r catalysis
butyl-3-me
BF4, liqui
trum of ap
Fig.(1
are the sal
eat interest
operties o
entirely
st importa
ductivity.(3
geous chass a uniqu
rous task-s
insufficie
philicity a
ironment.(
ride/coope
ids, espec
de or 1,3-
2), have a
widely us
ex ions a
s. The adv
ethylimida
id down to
pplications
-3):- Air an
lts having
ts recently
of ionic li
different
ant proper31,32)
aracterizue array o
specific ap
ently effec
and capabi(33)
8
er chlori
cially thos
-dialkylim
a broad ran
sed in elec
and electro
vent of a
azolium
o –81o C)
s involving
d water stab
g very low
y because
iquids hav
properties
ties of ele
zations ofof physico
pplication
ctive. Such
ility of pr
ide, 1/1,
se resultin
midazolium
nge of liqu
ctrochemic
onic abso
ir and wa
tetrafluoro
and (BM
g these liq
ble molten s
w melting
of their un
ve already
s from th
ectrolyte
f RTILs.ochemical
ns in which
h propertie
roviding w
, mp=25
ng from th
m chloride
uid phase
cal techno
orption sp
ater stable
oborate
MI.PF6) res
quids (Figu
alts.(28)
g temperat
nusual pro
y been m
hose of or
solutions
. propertie
h convent
es include
weekly coo
Introd
5o C).
he combin
e and alu
down to –
ologies as
ectra, for
e ionic liq
and hex
spectively,
ure 1-3).(2
ture. Ionic
operties as
mentioned
rdinary m
are non-v
es that ma
tional solv
e:
ordinating
duction
Organo-
nation of
uminium
–88 oC.
solvents
organic
quids, in
xafluoro-
, has led 29,30)
c liquids
s liquids.
in some
molecular
volatility
ake them
vents are
g or non-
Chapter One Introduction
9
2. Very good solvents properties for a wide variety of organic, inorganic and
organometallic compounds: in some cases, the solubility of certain solutes in
RTILs can be several orders of magnitude higher than that in traditional
solvents.(34)
3. Stability.
In most of these applications, the stability of ionic liquids, at least at a certain
extent, is crucial for optimum process performance. Several studies have indicated
that, although not 100% inert, certain ionic liquids incorporating 1,3-dialkyl
imidazolium cations are generally more resistant than traditional solvents under
certain harsh process conditions, such as those occurring in oxidation, photolysis
and radiation processes.(36)
4. Color.
High quality ionic liquids incorporating [bmim] + cation and a variety of anions,
such as [PF6]-, [BF4]-, [CF3SO3]-, [CF3CO2]- and [(CF3SO2)2N]- have been
reported to be colorless, even though they are not 100% pure. The color of less
pure ionic liquids generally ranges from yellowish to orange. The formation of the
color has been attributed to the use of raw materials with color or excessive heating
during the synthesis of imidazolium salt.(35,36)
5. Hygroscopicity.
The water content has an influence on the viscosity of the ionic liquids.
Viscosity measurements indicate that ionic liquids became less viscous with
increasing water content. Hydrolysis problems can also occur. Most of the ionic
liquids have extremely low vapour pressures, which allows removing water by
simple heating under vacuum. Water contents below 1 ppm are quite easy to
achieve with most of the liquids.
6. Hydrophopicity.
The degree of polarity can be varied by adapting the length of the 1-alkyl chain
(in 1,3-substituted imidazolium cations). The anion chemistry has a large influence
on the properties of IL, though little variation in properties might be expected
between same-cation salts of these species, the actual differences can be dramatic:
Chapter One Introduction
10
for example, [bmim]PF6- is immiscible with water, whereas [bmim]BF4
- is water-
soluble.(36)
7. Conductivity.
Ionic liquids have reasonably good ionic conductivities compared with those of
organic solvents/electrolyte systems (up to~10 mS cm-1). However, at room
temperature their conductivities are usually lower than those of concentrated
aqueous electrolytes. Based on the fact that ionic liquids are composed solely of
ions, it would be expected that ionic liquids have high conductivities. This is not
the case since the conductivity of any solution depends not only on the number of
charge carriers but also on their mobility(38). The large constituent ions of ionic
liquids reduce the ion mobility which, in turn, leads to lower conductivities. The
conductivity of ionic liquids is inversely linked to their viscosity. Hence, ionic
liquids of higher viscosity exhibit lower conductivity. Increasing the temperature
increases conductivity and lowers viscosity.(37,39)
8. Viscosity.
Generally, ionic liquids are more viscous than common molecular solvents and
their viscosities are ranging from 10 mPa s to about 500 mPa s at room
temperature. The viscosity of ionic liquids is determined by van der Waals forces
and hydrogen bonding(40). Electrostatic forces may also play an important role.
Alkyl chain lengthening in the cation leads to an increase in viscosity. This is due
to stronger van der Waals forces between cations leading to increase in the energy
required for molecular motion. Also, the ability of anions to form hydrogen
bonding has a pronounced effect on viscosity. The fluorinated anions such as BF4-
and PF6-form viscous ionic liquids due to the formation of hydrogen bonding, in
general, all ionic liquids show a significant decrease in viscosity as the temperature
increases.(41)
9. Density.
Ionic liquids in general are denser than water with values ranging from 1 to 1.6
g cm-3 and their densities decrease with increase in the length of the alkyl chain in
the cation. For example, in ionic liquids composed of substituted imidazolium
Chapter One Introduction
11
cations and CF3SO3- anion the density decreases from 1.39 g cm-3 for [EMIm]+ to
1.33 g cm-3 for [EEIm]+, to 1.29 g cm-3 for [BMIm]+ and to 1.27 g cm-3 for
[BEIm]+. The densities of ionic liquids are also affected by the identity of anions.
For example, the densities of 1-butyl-3 methylimidazolium type ionic liquids with
different anions, such as BF4-, PF6
-, TFA- and Tf2N- are 1.12 g cm-3, 1.21 g cm-3,
1.36 g cm-3 and 1.43 g cm-3, respectively.(42-44)
10. Melting point.
As a class, ionic liquids have been defined to have melting points below 100 oC
and most of them are liquid at room temperature. Both cations and anions
contribute to the low meting points of ionic liquids(43). The increase in anion size
leads to a decrease in melting point. For example, the melting points of 1-ethyl-3-
methylimidazolium type ionic liquids with different anions, such as [BF4]- and
[Tf2N]- are 15o C and -3o C, respectively. Cations size and symmetry make an
important impact on the melting points of ionic liquids. Large cations and
increased asymmetric substitution results in a melting point reduction.(45)
11. Thermal stability.
Ionic liquids can be thermally stable up to temperatures of 450o C. The thermal
stability of ionic liquids is limited by the strength of their heteroatom–carbon and
their heteroatom– hydrogen bonds, respectively. High temperatures are only
tolerated by most liquids for a short time. Long time exposure to such high
temperatures inevitably leads to decomposition.(46)
12. Electrochemical window.
The electrochemical window is an important property and plays a key role in
using ionic liquids in electrodeposition of metals and semiconductors. By
definition, the electrochemical window is the electrochemical potential range over
which the electrolyte is neither reduced nor oxidized at an electrode. This value
determines the electrochemical stability of solvents(47). As known in aqueous
solution, the electro-deposition of elements and compounds is limited by the low
electrochemical window of water (about 1.2 V). On the contrary, ionic liquids have
significantly larger electrochemical windows, e.g., 4.15 V for [BMIm] PF6 at a
Chapter One Introduction
12
platinum electrode, 4.10 V for [BMIm]BF4 and 5.5 V for [BMP]Tf2N at a glassy
carbon electrode. In general, the wide electrochemical windows of ionic liquids
have opened the door to electrodeposit metals and semiconductors at room
temperature which were formerly obtained only from high temperature molten
salts. For example, Al, Mg, Si, Ge, and rare earth elements can be obtained from
room temperature ionic liquids.(48)
1.7. Applications of Ionic liquids. 1. Homogeneous and heterogeneous catalyst.
For some applications, ionic liquids offer the advantage of both homogeneous
and heterogeneous catalysts. This is because selected ionic liquids can be
immiscible with the reactants and products but dissolve the catalysts. This has the
advantages of a solid for immobilizing the catalyst, with the advantages of a liquid
for allowing the catalyst to move freely.(49)
2. Biological reactions media; Enzymes are also stable in ionic liquids, opening
the possibility for ionic liquids to be used in biological reactions, such as the
synthesis of pharmaceuticals.(49,50)
3. Treatment of high-level nuclear waste; Ionizing radiation does not affect ionic
liquids, so they could even be used to treat high-level nuclear waste.(49)
4. Removing of metal ions; Researchers have designed and synthetized several
new ionic liquids to remove cadmium and mercury from contaminated water.
When these water-insoluble ionic liquids come in contact with contaminated water,
they snatch the metal ions out of water and sequester them in the ionic liquid.(51)
5. Purification of gases, Ionic liquids can selectively dissolve and remove gases
and could be used for air purification on submarines and spaceships.(51)
6. As solvent for synthetic and catalytic purposes, for example Diels-Alder
cycloaddition reactions, Friedel-Craft acylation and alkylation, hydrogenation and
oxidation reactions and Heck reactions.(52,55)
7. As biphasic system in combination with an organic solvent or water in
extraction and separation technologies. (52)
Chapte
8. As
easy rec
9. As
1.8. R
Am
The
behavio
form sta
quaterna
reasons.
and mol
vapor p
room-te
electroly
capacito
ammoni
when th
Fi
(I) M(II) M(III) M
er One
catalyst
cycling of
electrolyt
Room-Tem
monium
number
or is quite
able molt
ary ammo
. They hav
lten salt c
pressure fo
emperature
ytes for u
ors. Molt
ium salts,
he size of t
ig.(1-4):- Som
Molten Salts Molten Salts Molten Salts
immobiliz
homogen
tes in elec
mperatu
m Ion.
of familie
limited. E
en salts a
onium and
ve been us
chemistry,
or organic
e molten s
use in ba
ten salts
however
the cation
me types of
based on thbased on thbased on th
zation wit
eous catal
ctrochemis
ure Molt
es of mol
Excluding
at room te
d imidazo
sed as an
as well a
c reaction
salts have
atteries, p
have lon
r, the melt
and anion
room tempe
he imidazoliuhe bis(trifluohe quaternar
13
thout need
lysts. (52)
stry. (52)
ten Salt
lten salts
the hydrat
emperatur
lium salts
effective
as being e
ns and ins
e also been
photoelectr
ng been k
ting point
n (Figure 1
erature ioni
um cation. oromethylsury ammoniu
d of speci
ts Based
exhibitin
te salts, fe
e. Molten
s, have be
tool in ele
excellent i
strumental
n investig
rochemica
known in
t only app
1-4) is larg
ic liquid.
ulfonyl)imideum cation.
ial functio
d on th
g room te
ew purely
n organic
een invest
ectrolyte s
ionic med
l analysis
gated as p
al cells, e
the fam
proaches r
ge. (53)
e ion anion.
Introd
onalization
he Qua
emperatur
inorganic
salts, suc
tigated for
solution ch
dia with ne
. In recen
potential c
electroplat
ily of qu
room tem
duction
n for the
ternary
re liquid
systems
ch as the
r several
hemistry
egligible
nt years,
andidate
ting and
uaternary
mperature
Chapter One Introduction
14
1.9. General properties of Choline Chloride.
Choline chloride [Ethanaminium, 2-hydroxy-N, N, N-trimethyl-, chloride], is a
quaternary amine salt, with the chemical formula [HOCH2CH2N (CH3)3Cl] it
dissociates in water into the corresponding positively charged quaternary hydroxyl
alkylammonium ion and the negatively charged chloride ion. Using different
amounts of choline chloride showed that the lowest pH value of 4 was determined
in a100 mg/L. Even if choline chloride of different specifications was used, choline
chloride can be stated to be a weak acid. No data on dissociation constants are
available for this compound. Choline chloride has neither explosive nor oxidizing
properties due to its molecular structure (54). Choline is a dietary component and
found in foods as free choline and as esterified forms such as phosphocholine,
glycerol-phosphocholine, sphingomyeline, and phosphatidylcholine. Dietary
choline is absorbed from the lumen of the small intestine. Additionally to dietary
supply choline can be made available by enzymatic cleavage in the pancreas from
other nutritional sources, before choline can be absorbed from the gut, some is
metabolised by bacteria to form betaine and methylamines. Fasting plasma choline
concentrations vary from 9 to 20 μmol/L. Choline chloride is a white crystalline
solid but is marketed as an aqueous solution (70 – 75 % w/w in water) which is a
colorless liquid with an amine-like odor. It has a measured water solubility of ca.
650 g/L and a calculated vapor pressure of 6.57×10-10 Pa at 25°C.(55) Choline
chloride or trimethyl (2-hydroxyethyl) ammonium chloride is an essential feed
ingredient for optimum animal growth. Choline is classified as a vitamin and
belongs to the group of water soluble B-vitamins. Choline has three essential
metabolic functions :(56)
i) It is a key constituent of phosphatidylcholine and phosphocholine and
represents about 70 – 80 % of total phospholipids in the body.
ii) It plays a vital role in fat transport and metabolism, thus preventing
abnormal fat accumulation in the liver. It is also important in the establishment
and maintenance of cell membrane structure.
Chapter One Introduction
15
iii) Choline is a precursor to acetylcholine in impulse transmission of all nerves;
Choline is a source of free methyl groups, essential for numerous biological
functions. Animal studies with choline chloride show a low acute toxicity after
oral uptake. No acute toxicity attributable to choline was observed in humans
following oral doses of ≥3000 mg choline magnesium trisalicylate/day.
iv) Choline chloride does not show a mutagenic, clastogenic or DNA damaging
potential when tested in vitro; furthermore it has no structural alerts. There is
therefore no indication of a genotoxic potential in vivo.
1.10. Deep eutectic solvents.
Eutectic mixtures of salts have been utilized for a long time to decrease the
temperature for molten salt applications. In the extreme, ambient temperature
molten salts have been formed by mixing quaternary ammonium salts with metal
salts. This type of ionic liquid can be viewed as a deep eutectic resulting from the
formation of complex anions e.g. Al2Cl7 and Zn2Cl5 thus decreasing the lattice
energy and decreasing the freezing point of the system (57, 58). Work in this area
focused on chloroaluminate salts of imidazolium and pyridinium chloride. A
variety of different anions are formed in solution the ratios of which vary with the
changing aluminium chloride composition. The depression of freezing point at the
eutectic composition is as much as 190 °C. These ideas have recently been
extended to other chlorometallate salts including ZnCl2, SnCl2 and FeCl3.(59,60)
1.10.1. Eutectic mixtures of Urea and Choline chloride.
Halide salts can also form complexes with hydrogen bond donors and
previous work has shown that mixtures of urea with alkali metal halides form
eutectics with melting points of < 150° C. While a few reports also exist for
adducts of urea with other metal salts their use as solvents has been limited to high
temperature applications. Substituted quaternary ammonium salts such as choline
chloride with urea produce eutectics that are liquid at ambient temperature and
have unusual solvent properties. Figure (1-5) shows the freezing point of mixture
Chapter One Introduction
16
of choline chloride and urea, a eutectic occurs at a urea to choline chloride ratio of
2:1. The freezing point of the eutectic mixture is 12 °C, which is considerably
lower than that of either of the constituents (m.p. choline chloride = 302 °C and
urea = 133 °C) and allows the mixture to be used as an ambient temperature
solvent. This significant depression of the freezing point must arise from an
interaction between urea molecules and the chloride ion. This is consistent with the
crystallographic data for the (mane) solid [(CH3)3N+CH2CH2OH-]2C2O42-
.2(NH2)2CS which shows extensive hydrogen bonding between the thiourea
molecule and the oxalate anion. The existence of hydrogen bonding in these
eutectic mixtures can be observed using NMR spectroscopy. An alternative
approach to making large non-symmetric ions is to take a simple organic halide
salt and complex the anion with something that will form a hydrogen bond. The
complexing agent interacts with the anion, increasing its effective size. This in turn
decreases the interaction with the cation and so decreases the freezing point of the
mixture. The major advantage of this approach is that common, non-toxic
components can be used and the properties of the liquid can be changed with each
hydrogen bond formed and each halide salt. This approach has numerous
possibilities. Therefore, many quaternary ammonium salts are available and the
wide variety of amides, amines, carboxylic acids and alcohols which could be used
to provide hydrogen bonds. These liquids are similar to the metal-containing ionic liquids, but to differentiate them from the wide range of ionic liquids that are
available they have been given the name deep eutectic solvents. They are easy to
make – simply take the two solids, mix them together with gentle heating, until
they melt, and when they cool they remain liquid. Work to date has focused on
choline (2-hydroxyethyl-trimethylammonium) chloride as the quaternary
ammonium salt. This is vitamin B4 and is produced on the milirn tones per annum
scale as an additive for chicken feed and a wide range of other applications. Thus
choline chloride is not only cheap, but it is non-toxic and even biodegradable. Ionic
liquids based on choline chloride have, therefore, the potential to be used on a
large scale. One simple hydrogen bond ‘donor’ that has been used is urea
Chapte
((NH2)2
2:1 ratio
Fig.(1
This i
melts at
(termed
hydroge
compos
bonds, e
of 1:1 a
point an
providin
magnitu
and salt
ionic liq
the chlo
molten s
Thes
confirm
er One
C=O), wh
o produces
1-5):- Freezi
is remarka
t 303°C a
the ‘eute
en bond d
ition chan
e.g. oxalic
and a depr
nd the eu
ng the h
ude of the
. It was re
quid no p
oride is no
salt nature
se deep eu
ming that
hich is a co
s a colorle
ing point of
able given
and urea a
ectic com
donor is 1
nges with
c acid (HO
ression of
utectic co
hydrogen
freezing
eported th
recipitate
ot as free i
e.(59,76)
utectic sol
the ionic
ommon fe
ess liquid t
f Choline chl
n the melti
at 134° C)
mposition’)
1:2. The d
h the natu
OOCCOOH
f freezing
omposition
bonds. W
point dep
at with ad
was obse
ion in this
lvents are
species
17
ertiliser. U
that freeze
loride/Urea
ing points
). The low
) reported
depression
ure of the
H): cholin
point of 2
n changes
With dee
pression is
ddition of
erved for
s ionic liq
highly co
are disso
Urea mixed
es at ca 12
Mixture as
s of the co
west melti
d to occur
n of freez
e molecul
ne chloride
212°C.(61)
s with th
p eutecti
s over 10
sliver nitr
sliver chl
quid and th
onducting
ciated in
d with cho
2°C as in (
a function o
onstituents
ing point
r when th
zing point
le providi
e has a eut
The depre
he nature
ic solvent
times gre
rate to cho
loride, wh
he formed
g (ca. 1 m
the liqui
Introd
oline chlo
(Figure 1-5
of compositi
s (choline
of these m
he ratio o
t and the
ing the h
tectic com
ession of
of the m
ts, howev
eater than
oline chlor
hich indica
d ionic liq
mS cm-1 at
id and ca
duction
oride in a
5),
ion.(61)
chloride
mixtures
f salt to
eutectic
hydrogen
mposition
freezing
molecule
ver, the
with ice
ride/urea
ated that
quid is of
t 30 oC),
an move
Chapte
indepen
with tem
Fig. (1-6)function o
The p
applied
charge d
anion w
chloride
found to
ionic liq
While t
systems
mixture
Solvents
eutectic
water, m
are well
that hav
few repo
er One
ndently; (F
mperature.
:- Conductiof temperatu
principle
to mixtur
delocaliza
with the a
e with ure
o have inte
quids and
the depre
s, there is
s. To diff
s (DES) h
mixtures
many are b
l character
ve melting
orts also e
Figure 1-6
.
ivity and visure. (61)
of creatin
res of qua
ation is ac
amide mo
ea gave a
eresting so
a wide va
ssion of
a wide ra
ferentiate t
has been ad
s are easy
biodegrad
rized. The
g points of
exist for ad
6) shows
scosity of a
ng an ion
aternary a
chieved th
iety. It w
liquid wi
olvent pro
ariety of s
freezing
ange of am
these liqu
dopted. U
y to prepar
dable and t
e use of u
f <150 °C
dducts of u
18
that the
2 urea to 1
nic fluid b
ammonium
hrough hy
was shown
ith a freez
operties th
solutes we
point is
mides tha
uids from
nlike the r
re in a pu
the toxico
urea to for
has been
urea with
conductiv
choline chlo
by comple
m salts wi
ydrogen b
n that a e
zing point
hat are sim
ere found
less than
at can be u
ionic liqu
room-tem
ure state.
ological pr
rm eutectic
known fo
other met
vity incre
oride mole r
exing a h
ith a rang
bonding be
eutectic m
t of 12 °C
milar to am
to exhibit
n analogou
used to fo
uids the te
mperature i
They are
roperties o
cs with al
or a consid
tal salts.(61
Introd
eases sign
ratio ionic l
halide salt
ge of amid
etween th
mixture of
C. This liq
mbient tem
t high solu
us chloro
orm these
erm Deep
onic liquid
nonreact
of the com
lkali meta
derable tim, 63)
duction
nificantly
liquid as a
t can be
des. The
he halide
f choline
quid was
mperature
ubilities.
metalate
eutectic
Eutectic
ds, these
tive with
mponents
l halides
me and a
Chapter One Introduction
19
1.10.2. Eutectic mixtures of Malonic acid and Choline chloride.(83)
Carboxylic acids can also form complexes with hydrogen bond donors and
previous work has shown that mixtures of urea with quaternary ammonium salt
(choline chloride) form eutectics with melting points of 12° C.
Another hydrogen bond ‘donor’ that has been used is malonic acid
[CH2 (COOH)2], which is dicarboxylic acid. Malonic acid was mixed with choline
chloride in mole ratio of 1:1 and was heated to 80oC, when a homogeneous
colourless liquid was produced. The liquid was slowly cooled until it solidified at
0oC; freezing point was taken as the temperature at which the first solid began to
form. The solidified melt turned into liquid state as temperature increased to room
temperature which is considerably lower than that of either of the constituents
(choline chloride melts at 303°C and malonic acid at 133° C). The bonding was
suggesting being between the acid and chloride ion, or bridging acids between
neighboring chloride ions. (62) Phase Behaviour. Figure (1-7) shows the phase diagram for mixtures of choline
chloride with phenylacetic and phenylpropionic acids as a function of composition.
It is apparent that a eutectic is formed at a composition of 67 mol % acid which is
similar to the eutectic recently reported for the choline chloride/urea system. It can
be inferred that to form the eutectic two carboxylic acid molecules are required to
complex each chloride ion. It is thought that the protons and carboxylate anions are
extensively associated in the liquids.(62) (Figure 1-8) shows the phase diagrams for
dicarboxylic acid (malonic acid). For these systems, the eutectic occurs at 50 mol
% acid suggesting a 1:1 complex between the acid and chloride ion, or bridging
acids between neighboring chloride ions, like malonic acid. These exhibit the
rheology of gels and presumably have extensive bridging of the acids between
neighboring chloride ions. The freezing point of the eutectic mixture of choline
chloride and malonic acid mixture is 10oC, while the melting point of the pure acid
is 135oC and for choline chloride is 303oC.
Chapte
Fig.(1-7):function o
The tr
in the re
have ex
lowest m
malonic
molecul
Fig.(1-8):function o
The r
although
Mixture
choline
er One
- Freezing pof compositi
riscarboxy
egion 30-3
xtensive b
melting po
c acid is 1
le that don
- Freezing of compositi
resultant m
h they ar
es of zinc
chloride h
points of Cion.(62)
ylic acids c
35 mol %
bridging o
oint of thi
:1. The ch
nating the
points of Cion.(62)
molten sa
re, in gen
c chloride
have large
Choline chlor
citric and
acid. The
of the aci
is eutectic
hanging in
hydrogen
Choline chl
lts have th
neral, mo
with sub
e depressio
20
ride with p
tricarbally
se exhibit
ids betwe
c occurs w
n freezing
n bonding.(
loride with
he advant
ore viscou
bstituted q
ons of free
henylpropio
ylic acid h
t the rheol
en neighb
when the
point dep(83)
Oxalic, M
age that th
us than t
quaternary
ezing poin
onic and ph
have a eut
ogy of gel
bouring ch
ratio of ch
pended on
alonic, and
hey are no
their alum
y ammoni
nt (up to 2
Introd
henylacetic
tectic com
ls and pre
hloride io
holine chl
n the natur
Succinic a
ot water s
minum an
ium salts
70 °C). It
duction
acids as a
mposition
sumably
ons. The
loride to
re on the
acids as a
sensitive,
nalogues.
such as
has also
Chapter One Introduction
21
recently been shown that these ideas can be extended to other metals by using
hydrated salt mixtures. To circumvent the use of metals in ionic fluids a variety of
closed shell anions have been synthesized and these have been the main focus of
ionic liquids research over the past decade.
The viscosity of the DES of choline chloride/malonic acid changes significantly
as a function of temperature, the hydrogen bond donor type and mixture
composition as shown in (Figure 1-9). The viscosity values were found to cover the
range 50 to 5000 cP, which is similar to those observed with ionic liquids.
The DES of choline chloride/malonic acid has conductivity of the range 0.1 to
10 mS cm-1 changing with composition and temperature. These values are again
similar to imidazolium based ionic liquids and similar to the choline chloride/urea
DES. Approximately linear correlations are observed between σ and η-1 for this
DES for all of the systems, although the slopes are markedly different.(62)
Fig.(1-9):- Plot of log viscosity vs. reciprocal of temperature for the Choline chloride/Malonic acid ionic liquid. (62)
1.11. Interaction of the ligands with the metal ions. The tendency of metal ion to from a stable complex with ligands depend on
many rules such as the hard and soft acids and bases (HSAB) rule of Pearson (58)
which imply that metal ion tend to coordinate with certain functional groups of the
ligand to from a stable complex. On the other hand, the tendency of transition
Chapter One Introduction
22
metal ion of a special oxidation states is affected by the coordination to certain
ligands. This phenomenon is called (symbiosis).(64, 65)
Increasing the positive charge on the central transition metal ions strengthens
the metal –ligands bond. The metal ion prefers to bind with atoms of high electron
density such as N3-, O2-, P3-, S2- and C4-. The ligand should have certain
characteristic properties to make it convenient to form stable complex with
transition metal ions. The size, geometrical shape, number and geometrical
arrangement of ligand and donor atoms play the important role in the stability of
the resultant complex. Metal centers, being positively charged, are favored to bind
to negative charged biomolecules, the constituents of proteins and nucleic acid
offer excellent ligands for binding to metal ions (65). The Irving Williams series of
stability for a given ligand shows a good criterian for the stability of complexes
with dipositive metal ions which follows the order:
Ba2+<Sr2+<Ca2+<Mg2+<Mn2+<Fe2+<Co2+<Ni2+<Cu2+>Zn2+
This order arises in part from a decrease in size across the series and in part
from ligand field effect. A second observation is that certain ligands form their
most stable complexes with metal ions such as Ag+, Hg2+and Pt2+, but other ligands
seem to prefer ions such as Al3+, Ti4+, and Co3+. Ligands and metal ions were
classified as class (a) or (b) according to their preferential bonding. Class (a) metal
ions includes those of alkali metals, alkaline earth metals, and lighter transition
metals in higher oxidation states such as Ti+4 ,Cr+3 , Fe+3 ,Co+3 and the hydrogen
ion, H+. Class (b) metal ions include those of the heavier transition metals, and
those in lower oxidation states such as Cu+, Ag+, Hg+2, Pd+2, and Pt+2. According
to their preference toward either class (a) or class (b) metal ions, ligands may be
classified as type (a) or (b), respectively. Stability of these complexes may be
summarized as follows (65, 66); the ligand should have certain characteristic
properties to make it convenient to form a stable complex with transition metal
ions. The size, geometrical shape, number and geometrical arrangement of ligand
donor atoms play the important role in stability of the resultant complex (65).
Chapter One Introduction
23
Tendency of ligands to complex with Tendency of ligands to complex with
Class (a) metal ions Class (b) metal ions N>>P> As >Ab N<<P< As >Sb O>>S>Se>Te O<<S<Se~Te F>Cl>Br>I F<Cl<Br<I
One of the important aspects of CFT is that all ligands are not identical when
it comes to a causing separation of the d-orbitals. For transition metal compounds,
they were aware of the multitude of colors available for a given metal ion when the
ligands or stereochemistry are varied. In octahedral complexes, this can be
considered a reflection of the energy difference between the higher dz2, dx2-y2 (eg
subset) and the dxy, dyz, dxz (t2g subset).
It has been established that the ability of ligands to cause a large splitting of the
energy between the orbitals is essentially independent of the metal ion and the
spectrochemical series is a list of ligands ranked in order of their ability to cause
large orbital separations.
A shortened list includes:
I- < Br- < SCN- ~ Cl- < F- < OH- ~ ONO- < C2O4 2- < H2O < NCS- < EDTA4- < NH3
~ pyr ~ en < bipy < phen < CN- ~ CO
When metal ions that have between 4 and 7 electrons in the d orbitals form
octahedral compounds, two possible electron allocations can occur. These are
referred to as both weak field (high spin) and strong field (low spin)
configurations.
The MOT method has provided a quantitative measure of the d-splittings. The
eg/t2g splitting (called the octahedral ligand field splitting ∆0= 10Dq) Follows the
general trend. Point to consider about the ∆0 splitting:
1) charge on the metal:-
As charge increase, electrostatic attraction between M and L increases, M-L
bond distance decreased and SML increased. Also as charge increase, metal
Chapter One Introduction
24
becomes more electronegative, decreasing ∆EML. Both trends lead to greater
field strength for the more highly charged ion.
2) The nature of the metal ion, radial extension of 2nd
and 3rd
row transition metals
are greater, thus SML larger, leading to the following trend in ∆0: 1st
row
transition metal << 2nd
row transition metal ~ 3rd
row transition metal.
3) Nature of ligand:-
Different ligands have different SML for a given metal ion. Can assess by measuring ∆0 for different ligands about a given metal ion of given charge
Cl -< F
-< OH2 < CN
- Cl
-< OH2 < NH3 < CN
-
From experiments such as these, may deduce a general ranking of ligands in terms
of field strength… this ranking is called the spectrochemical series:
I- < Br
- < S
2- < SCN
- < Cl
-< S
2- < F
- < O
2- < OH
-< H2O < NH3 < CN
- < CO < NO
+
Weak field strength (small ∆0) strong field strength (large∆0) π-donors σ-only π-accepter
In the above series, the ligand type has been overlayed onto the spectrochemical
series. As is readily apparent from the energetic disposition of ∆0, π-donors give
weak ligand fields, σ-only ligands give intermediate fields and π-acceptors ligands
give strong fields. (64, 65)
Chapter One Introduction
25
1.12. The Chemistry of the first Transition Metal Elements.
1.12.1. Chromium (III), (d3).
Oxidation state (III) is by far the most stable and best known oxidation state of
chromium there is literally thousands of chromium (III) complexes that with a few
exceptions are all hexacoordinate. The principale characteristic of these complexes
in aqueous solution is their relative kinetic inertness.(64)
Chromium (III) forms stable salt with all the common anions and it complexes
with virtually any species capable of donating an electron-pair. These complexes
may be anionic, cationic, or neutral and, with hardly any exceptions are octahedral.
There is also a multitude of complexes with two or more different ligands such as
pentammines [Cr(NH3)5X]n+1 which have been extensively used in kinetic. These
various complexes are notable for their kinetic inertness, which is compatible with
the half-filled t2g level arising from an octahedral d3 configuration, ligand
displacement or rearrangement reaction of Cr (III) complexes are only about 10
times faster than those of Co(III), with half-times in the range of several hours.
This is the reason why many thermodynamically unstable complexes of Cr III can
be isolated as solid and that they persist for relatively long of time in solution, even
under conditions of marked thermodynamic instability. (65)
[CrX6]3- (X= halide, CN, SCN, N3, H2O, NH3).
[Cr(L-L)3] (L-L= amino acid anions).
The hexaaqua ion [Cr(H2O)6]3+, which is regular octahedral, occurs in aqueous
solution and in numerous solid salts such as the violet [Cr(H2O)6]Cl3. Chromium
(III) is the archetypal d3 ion and the electronic spectra and magnetic properties of
its complexes have been exhaustively studied. The magnetic properties of the
octahedral Cr(III) complexes are uncomplicated. All such complexes must have
three unpaired electrons irrespective of the strength of the ligand field, and this has
been confirmed for all known mononuclear complexes. The spectra of Cr(III)
complexes are also well understood in their main features, in an octahedral field,
the splitting of free ion ground 4F term, along with the presence of the excited 4P
Chapter One Introduction
26
term of the same multiplicity, provides or indicating the possibility of three spin-
allowed d-d transitions, and these have been observed in a considerable number of
complexes, the spectrochemical series was originally established by Tsuchida
using data for Cr(III) and Co(III) complexes. In the aqua ion the bands are found to
be 17,400, 24,700, and 37,000 cm-1. The frequencies of the spin-allowed bands of
Cr(III) in ruby indicate that the Cr(III) ions are under considerable compression,
since the value of Δo calculated is significantly higher than the [Cr(H2O)6]3+ion or
in other oxide lattice and glasses. Spin-forbidden transitions from the 4A2g ground
state to the doublet states arising from 2G state of the free ion are observed. The
transitions to the 2Eg and 2T1g states give rise to extremely sharp lines because the
slope of the energy lines for these states is the same as that for the ground state.
For [Cr(H2O)6]3+, this has the simplifying effect that the crystal field splitting
10Dq, is given directly by the energy of the lowest transition. The magnetic
moment is expected to be very close value to the spin-only value for three unpaired
electrons, and because of the absence of any orbital contribution, to be independent
of temperature. (65)
In molten salts anhydrous Chromium (III) chloride dissolved only slowly
forming pale green solutions after more than 15 h for 10−3 mol dm−3 concentrations
at 90oC in acetamide-potassium nitrate and 8 h at 140oC in butyramide-sodium
nitrate melts. The spectrum for the acetamide eutectics (Figure 1-10, curve A)
showed two absorption maxima, consistent with octahedral coordination. The band
energies were close to those expected for 6 acetamide or 6 nitrate ligands, but the
much higher extinction coefficients suggested mixed, lower symmetry, complexes,
though probably with little or no coordination by chloride ligands. The
coordination in the butyramide-sodium nitrate eutectic, spectrum in (Figure1-10,
curve B), was probably rather similar.(67)
Chapte
Fig.(1-10)(•), curve 1.12.2.
Iron (
with it b
one of th
to hydro
(equival
initial st
because
into the
noncom
The m
pyramid
very low
like ED
ligands
polyols
complex
affinity
much le
er One
):- Spectra oB - c = 0.00
Iron (II
(III) occu
because o
he most c
olyze and
lent in the
tage. The
e of charge
e visible
mplexing an
majority
dal ones a
w. No sim
DTA form
that coor
such as
x [Fe(C2O
for F- is
ess is kno
of CrCl3, cu03 mol dm−3
II), (d5).
urs in salts
of their ch
onspicuou
d/ or form
e first stag
various h
e transfer
region.
nion are y
of Iron c
are also im
mple ammi
some def
rdinate by
glycerol
O4)3]3-. Fer
quite hig
own of th
urve A - c = in butyram
s with mo
haracter as
us features
complexe
ge to acid
hydroxo s
band in t
Thus aq
yellow unl
complexes
mportant.
ine compl
finite com
y oxygen
and suga
rric ion fo
gh. Iron (
he details
27
0.006 mol dide-sodium
ost anions
s reducing
s of ferric
es. It has
dissociatio
species, su
the ultravi
queous so
ess strong
s are octa
The affin
lex exists
mplexes. I
n, viz., ph
ars, with
orms comp
(III) is iso
of FeIII s
dm−3 in acetnitrate at 14
s, except t
g agent. In
iron in aq
been esta
on of the
uch as [Fe
iolet regio
olutions o
gly acid.(65
ahedral, b
nity of Iro
in aqueou
Iron (III)
hosphate
oxalate I
plexes wit
oelectronic
spectra be
tamide-pota40oC (- - -).(6
those that
n aqueous
queous sol
ablisned th
aqua ion)
e(OH)(H2O
on, which
of ferric 5)
but tetrah
on (III) for
us solution
has its gr
ions, pol
ron (III)
th halide i
c with ma
ecause of
Introd
assium nitra67)
t are incom
s solution
lution its t
hat the hy
is govern
O)5]2+, are
have tails
salts ev
hedral and
r amine li
n, chelatin
reatest aff
lyphospha
forms tri
ions and S
anganese
the much
duction
ate at 90oC
mpatible
Fe (III)
tendency
ydrolysis
ned in its
e yellow
s coming
en with
d square
igands is
ng amine
finity for
ates, and
isoxalate
SCN-, its
(II), but
h greater
Chapter One Introduction
28
tendency of the trivalent ion to have charge transfer bands in the near-ultraviolet
region with strong low-energy wings in the visible that obscure the very weak,
spin-forbidden d-d band. Insofar as they are known, however, the spectral features
of iron (III) in octahedral surroundings are in accord with theoretical expectations.
Iron (III) is high spin in nearly all its complexes except those with the strongest
ligands like [Fe (CN)6]3-, [Fe (bipy)3]3+, [Fe (phen)3]3+, and other tris compexes
with imine nitrogen atoms as donor. The low-spin complexes with t2g5
configuration usually have considerable orbital contribution to their moments at
about room temperature. There is evidence of very high covalence and electron
delocalization in low-spin complexes such as [Fe (phen)3]3+ and [Fe (bipy)3]3+. Iron
(III) may be high or low spin and even though an octahedral stereochemistry is the
most common. A number of other geometries are also found. It has been suggested
that FeIII species high spin and low spin are present in thermal equilibrium. The
simplified Tanabe-Sugano diagram (d5) shows how as the crystal field increases, it
is designated as weak or strong (the configuration as high or low spin) depending
on whether the ground term is a 6A1 (t2g3eg2) or 2T2 (t2g5) term. Where CFT is close
to this critical boundary value both terms are thermally accessible and the observed
magnetic moment will depend on the precise distribution between them. As this
distribution will be temperature dependent. (64, 66)
Iron (III) chloride was very soluble in the acetamide-potassium nitrate eutectic
at 90oC giving initially a pale yellow solution which darkened rapidly to a strong
brown color within 15 min. The pale yellow solution gave one absorption band at ν
= 16900 cm−1 (ε = 35mol−1 dm3 cm−1) close to an absorption edge at ν = 21000
cm−1, while the dark brown solution showed only the absorption edge. The pale
yellow solution may be due to a tetrahedral complex perhaps containing some
chloride (FeCl4- in LiNO3-KNO3 eutectic at 160oC, band at ν = 16000 cm−1, ε =
32mol−1 dm3 cm−1) which was replaced over 15 min by more stable simple
complexes containing acetamide and/or nitrate, or perhaps by more polymerized
solvolyzed iron specie.(67)
Chapter One Introduction
29
1.12.3. Cobalt (II), (d7).
There is a very marked contrast in this oxidation state between cobalt (II) on the
one hand and the two heavier member of the group on the other hand, for cobalt
(II) it is one of the two most stable oxidation state. Complexes of configuration
comparable are stability with the t2g6 of Co (II), show a greater diversity of types
and are more labile. The aqua ion [Co (H2O)6]+2 is the simplest complex of cobalt
(II). In aqueous solution containing no complexing agents, the oxidation to CoIII is
very unfavorable. Cobalt (II) forms numerous complexes, mostly either octahedral
or tetrahedral but five-coordination and square species are also known. There are
more tetrahedral complexes of Cobalt (II) than other transition metal ions. This is
in accord with the fact that for a d7 ion, ligand field stabilization energies disfavor
the tetrahedral configuration relative to the octahedral one to a smaller extent than
for any other dn (1 ≤ n ≤ 9) configuration. Because of the small stability difference
between octahedral and tetrahedral CoII complexes, there are several cases in which
the tow types with the same ligand are both known and may be in equilibrium.
There is always some [Co (H2O)4]2+ in equilibrium with [Co(H2O)6]2+.(64) As
already noted, Cobalt (II) occurs in a great variety of structural environments,
because of this the electronic structures, hence the spectral and magnetic properties
of the ion, are extremely varied. High-spin octahedral and tetrahedral complexes,
in each case there is a quartet ground state (4F) and three spin-allowed electronic
transitions to the excited quartet state (4P). As maybe inferred from the simple
observation that octahedral complexes are typically pale red or purple and many
common tetrahedral ones are an intense blue in each case the visible spectrum is
dominated by the highest energy transition, 4A2→4T1(P) for tetrahedral and 4T1g(F)→4T1g(P) for octahedral complexes, however in the octahedral systems the 4A2g level is usually close to the 4T1g(P) level and the transition to these tow levels
are close together. Since the 4T1g state is derived from a t2g3eg4 electron
configuration, and the 4T1g (F) ground state is derived mainly from a t2g5eg2
configuration, the 4T1g(F)→4A2g(F) (shoulder) transition is essentially a two-
Chapter One Introduction
30
electron process. For octahedral complexes, there is one more spin-allowed
transition (4T1g (F) →4T2g) which generally occurs in the near-infrared region.
Tetrahedral complexes [CoX4]2- are generally formed with monodentate anionic
ligands such as Cl-, Br-, I-, SCN-, N3-, and OH-; with a combination of two such
ligands and two neutral ones, tetrahedral complexes of the type CoL2X2 are
formed. With ligands that are bidentate monoanions, tetrahedral complexes are
formed. For tetrahedral complexes there is also a transition in the near-infrared
region [4A2→4T1(F)] as well as one quite low energy (4A2→4T2), which is observed
because it is in region spectrum(1000-2000 nm)and it is orbitally forbidden. The
visible transitions in both tetrahedral and octahedral cases, particularly in the
tetrahedral case is more intense, because a number of transitions to doublet excited
state occur in the same region, and these acquire some intensity by means of spin
orbit coupling. Cobalt (II) is the only common d7 ion, because of its stereochemical
diversity its spectra have been widely studies. In the octahedral case the splitting is
the same as the octahedral d2 ion and the spectra can therefore be interpreted by
using the same Tanabe-Sugano diagram as was used for d2. Magnetic properties
provide a complementary means of distinguishing stereochemistry the T ground
term of octahedral ion is expected to give rise to a temperature dependent orbital
contribution to the magnetic moment where as the A ground term of the tetrahedral
ion is not, Low-spin octahedral complexes, a sufficiently strong ligand field (Δo ≥
15,000 cm-1) can cause a 2E state originating in the 2G state of free ion to become
the ground state. The electron configuration here is mainly t2g6eg; thus a Jahn-
teller distortion would be expected consequently perfectly octahedral low-spin Co
(II) complexes must be rare.(64,65)
The dipositive 3d transition metal ions have been shown to occur as markedly
different ionic species in the highly dissociated alkali chloride melts compared
with the non dissociated aluminum chloride melt. From spectrophotometric
measurements of Cobalt (II) chloride in calcium nitrate tetrahydrate melts
containing variable amounts of calcium chloride (from 0 to 3.1 mol dm-3) were
reported at different temperatures from 40 oC to 70oC. The solutions in pure
Chapter One Introduction
31
calcium nitrate tetrahydrate showed octahedral coordination, the position of the
absorption maximum being at 19420cm-1 (512 nm). Addition of chloride caused a
shift of the absorption maximum toward lower energies and a large increase of the
absorption coefficient indicating a change from octahedral to tetrahedral or
distorted octahedral coordination.(68)
The absorption spectra of Cobalt (II) chloride complexes, containing variable
concentrations of chloride ligand also reported, in a molten mixture of 80% acetmide-
20% calcium nitrate tetrahydrate, in range of temperature 313 to 363 K, in the
wavelength range 400-800 nm. The melt contained three ligands (CH3CONH3, H2O,
and NO3-) for competition with added chloride; addition of chloride caused a shift of
the absorption maximum of octahedral Cobalt (II) nitrate towards lower energies and
pronounced changes in shape of the initial spectrum of Cobalt (II) nitrate. The effect
of temperature changes on the molar absorption of the Co (II) species was attributed
to structural changes occurring in the Cobalt (II) species.(69)
In the LiC1-KC1 eutectic melt the absorption spectrum of Co (I1) is that of the
tetrahedral CoCl42- species. The occurrence of the CoCl4
2- rather than the CoCl64-
species in this highly ionic melt can be rationalized in terms of large (-30 kcal.)
electrostatic repulsion energies favoring the four- over the six-coordinated form. In
the covalent A1C13 melt CoCl2 in an A1C13 melt with 35.5 mole % KC1, close to the
composition KA12C17, is practically identical with the spectrum in pure A1C13,
indicating that the postulated Co(Al2C17)2 complex is the only light-absorbing species
present in the composition range 0-35.5 mole % KC1. Successive additions of KC1 to
the 35.5 mole % KC1 melt resulted in only minor spectral changes until a
composition of 42 mole % KC1 was reached. Upon further addition of KC1, the
Co(Al2C17)2 spectrum characteristic of octahedral coordination gradually transformed
into a new spectrum. Addition of KC1 to a melt containing 49.3 mole % KC1
resulted in partial precipitation of Co (I1) and the formation of a blue precipitate
which chemical and X-ray analysis showed to be CoCl2. Addition of KC1 to a melt
whose KC1:AlCl3 ratio was exactly equal to 1 caused the CoCl2 precipitate to
redissolve completely with only a few times more than the stoichiometric amount of
KC1 needed to form CoC142-, as shown in (Figure 1-11).(70)
Chapte
Fig.(1-11):49.7%AlCCo(Al2Cl7)distorted o Absor
room-tem
butylpyr
(I1) was
Spectros
octahedr
suggeste
molar ra
Fig. (1-12Co (II) inmelt (tetr
er One
:- Spectra ofCl3 (tetrahedr)2 ; C, 100.0 octahedral; E
rption spe
mperature
ridinium c
s tetrahedra
scopic stud
rally coor
ed that two
atio compo
2):- Absorptn 2:l AICI3,rahedral).(71)
f Co(II) in cral CoCl4
2-), Emole % KC
E, 35.5 mole %
ctroscopy
chloroa
chloride (A
ally coord
dies of Co
rdinated sp
o different
osition regi
tion spectra -BPC melt )
chloride surrE, 50.3 mole
C1 at 800o; D% KCl, 64.5 m
was used
aluminate
AICI3-BPC
dinated as t
o (I1) in 2
pecies, po
t Co(I1) s
ion. As sh
of Co (II) i(octahedral
32
roundings: Ae % KC1, 49D, 49.9 mole mole % AlC1
d to study
melt
C). Absorp
the CoCI42
2: 1 molar
ossibly Co
pecies ma
own in (Fi
in AICI3,-BPl); (dotted li
A, crystalline9.7 mole % A% KC1, 50.1
13 at 300o. (70)
the coordi
system
rption spec2- ion in 0.
r ratio A1
o (Al2C I7
ay exist in
igure1-12)
PC melt at 4ine) 1.02 ×1
e Co(AlC14)2
AlCl3 at 300o
1 mole % M
ination of
aluminum
ctroscopy
.8: 1 molar
CI3-BPC
7) m2-m , w
equilibriu
).(71)
40.0 OC: (so10-3M Co(II)
Introd
2 at 25o, B,5o (octahedral
MCl2 at 300o[C
Co (I1) io
m chlor
indicated
r ratio AIC
melt sugg
with m =
um close t
olid line) 2.0) in 0.8:l AI
duction
0.3% KCl, CoCl6
4- or Co(AlCl4)2]
on in the
ride-N-n-
that Co
CI3-BPC.
gested an
= 2, also
to the 1:l
04 ×10-2 M ICI3,-BPC
Chapte
In ano
butyram
The visi
showed
coordina
Environ
acetami
graduall
while th
change w
nitrates
tetrahed
melts w
coefficie
an even
strain im
Fig. (1-1temperatucurve C (
er One
other stud
mide-sodiu
ible- ultra
two band
ation in
nment sug
de-potassi
ly increase
he band i
was in acc
replacing
dral coordi
with increa
ents found
n lower sy
mposed by
3):- Spectrures, curve o).(67)
dy anhydro
um nitrate
aviolet spe
ds, with s
each cas
ggest coo
ium nitrat
ed to 140o
intensities
cord with
g acetam
ination, as
ased tempe
d with but
ymmetry c
y the large
ra of CoClA 90oC (•)
ous cobalt
and acetam
ectrum of
ome splitt
se. Calcul
ordination
te eutecticoC, the ab
s increase
a shift to
ide or p
s has frequ
erature. Th
tyramide-s
coordinati
r butyram
l2 (0.019 m), curve B 14
33
t (II) chlo
mide-pota
the acetam
ting of th
lations u
by 4 a
c. As the
bsorption b
d conside
a lower sy
erhaps a
uently bee
he larger
sodium nit
ion, possib
mide molec
mol dm−3) 40oC (- - -),
oride was
assium nitr
mide solut
e visible b
sing Jorg
acetamides
temperat
bands mov
erably (Fi
ymmetry e
shift fro
en found f
red shifts
trate (Figu
bly due a
cules.(67)
in acetamiand of buty
found ver
rate, givin
tion (Figu
band, ind
gensen's R
s and 2
ture of the
ved a littl
igure 1-13
environme
om octah
for cobalt
and even
ure 1-13, c
at least in
ide-KNO3 yramide-sod
Introd
ry soluble
ng violet so
ure 1-13, c
dicating oc
Rule of
chlorides
ese soluti
e towards
3, curve B
ent, possib
hedral to
(II) salts
n higher ex
curve C) i
part to th
eutectic at dium nitrate
duction
e in both
olutions.
curve A)
ctahedral
Average
s in the
ions was
s the red,
B). This
bly some
a more
in many
xtinction
ndicated
he steric
different e at 140oC,
Chapter One Introduction
34
1.12.4. Nickel (II), (d8). This is undoubtedly the most prolific oxidation state for this group of elements.
The stereochemistry of Ni (II) has been atopic of continuing interest, and kinetic and
mechanistic studies on complexes of Ni (II). The absence of any other oxidation state
of comparable stability for nickel implies compounds of Ni (II) are largely immune to
normal redox reactions. Ni (II) forms salts with virtually every anion has extensive
aqueous solutions chemistry based on the green hexaaaquanickel(II), [Ni(H2O)6]2+ ion
which is always present in the absence of strongly complexing ligands. The
coordination number of Ni(II) rarely exceeds 6 and its principal stereochemistries are
octahedral and square planer, tetrahedral (4-coordinate) with rather fewer examples of
trigonal bipyramidal, square pyramidal (5-coordinate). Octahedral complexes of Ni
(II) are obtained often from aqueous solution by replacement of coordinating water,
especially with neutral N-donner ligands such as NH3, en, bipy, and phen, but also
with NCS-, NO2-, and O-donner such as DMSO (Me2SO). (64, 65)
Nickel (II) forms a large numbers of complexes encompassing coordination
numbers 4, 5, and 6 in structural types of neutral ligands. It is characteristic of
Ni(II)complexes that complicated equilibria, which are generally temperature
dependent and sometimes concentration dependent. Nickel(II) is the only common d8
and its spectroscopic and magnetic properties have accordingly been extensively
studies in cubic field three spin-allowed transitions are expected, because of the
splitting of the free ion ground 3F term and the presence of the 3P term, in an
octahedral field the splitting is the same as for the octahedral d3 ion and the same
Tanaba-Sugano diagram, can be used to [Ni(H2O)6]2+, the spectra can readily be
interpreted by referring to the energy level diagram for d8 ions. Three spin-allowed
transitions are expected, and the three observed bands in each spectrum may thus be
assigned. It is a characteristic feature of the spectra of octahedral nickel (II)
complexes like [Ni(H2O)6]2+, the molar absorbances of the bands are at low end of the
range (1-100) for octahedral complexes of the first transition series in general, the
splitting of the middle band in the [Ni(H2O)6]2+ spectrum is due to spin-orbit coupling
that mixes the 3T1g(F) and 1Eg states, which are very close in energy at the Δo value
given by 6 H2O. The assigned bands are, ν1 (3A2g (F) → 3T2g (F)) is the 10Dq it is
Chapte
also evid
absorpti
spin-orb
spin-forb
(3A2g(F)
this is fo
there are
of a p
specifica
Anhy
solution
indicatin
slightly
coordina
acetamid
complex
no data
similar t
and 2 ch
Fig. (1-14(•), curve
er One
dence of w
on has a s
bit couplin
bidden tra
)→3T1g(P)
ound alway
e character
paramagne
ation.(64,65)
ydrous N
s. The spe
ng octahed
lower ene
ation by c
des and 2
xes resultin
on the b
to that of
hloride liga
4):- Spectra B - c = 0.01
weak spin-
strong sho
g ‘mixing
ansition t
)), it follow
ys to be th
ristic spec
etic nicke)
ickel (II)
ectra (Figu
dral coordi
ergies than
chloride, o
chlorides
ng would e
and positi
acetamide
ands.(67)
of NiCl2, cu2 mol dm−3
-forbidden
oulder on
’ (spin sin
to gain in
ws that th
he case, for
ctral and m
el (II) c
chloride
re 1-14, cu
ination of
n those fou
or less lik
could giv
explain the
ions of bu
e, then the
urve A - c = in butyram
35
n (spin trip
it. This ha
nglet (1Eg)
ntensity f
hey all sho
r regular o
magnetic pr
complex t
was solub
urves A an
the Nicke
und in pu
kely by n
ve the obs
e higher e
utyramide
calculated
0.046 mol dide-sodium
plet to sing
as been a
) with the
from the
ould have
or nearly re
roperties,
the less
ble in bo
nd B) show
el (II) catio
re molten
nitrate. Jor
erved ban
xtinction c
complexe
d coordina
dm−3 in acetnitrate at 14
glet), ν2 (3
scribed to
(3T1g) the
spin-allow
tow unpai
egular tetr
more irreg
likely c
oth eutecti
wed three a
ons. The b
acetamide
rgensen's
nds and the
coefficient
es, but if
ation woul
tamide-pota40oC (- - -).(6
Introd
3A2g (F)→
o the influ
ere by allo
wed trans
ired electr
rahedral co
gular the g
onform t
ics, giving
absorption
bands are s
e suggesti
Rule sugg
e less sym
ts found. T
its ligand
ld also be
assium nitra67)
duction
→3T1g(F))
uences of
wing the
ition, ν3
rons, and
omplexes
geometry
to these
g yellow
n maxima
shifted to
ing some
gested 4
mmetrical
There are
d field is
4 amide
ate at 90oC
Chapter One Introduction
36
1.12.5. Cupper (II), (d9).
The dispositive state is the most important one of Copper. In this oxidation state
Copper which provide by for the most familiar and extensive chemistry. Simple salt
are formed with most anions, except CN- and I-. Which instead form covalent Cu(I)
compounds which insoluble in water. Most Cu(I) compounds are fairly readily
oxidized to Cu(II) compounds. there is a well-defined aqueous chemistry of Cu2+, and
large number of salts of various anions, many of which are water soluble, the blue
color of their solution arising from the [Cu(H2O)6]2+ion, and they frequently
crystallize as hydrates. The aqueous solutions are prone to slight hydrolysis and,
unless stabilized by a small amount of acid, are liable to deposit basic salts. The most
common coordination numbers of Copper (II) are 4, 5, and 6, but regular geometries
are rare and the distinction between square-planar and tetragonally distorted
octahedral coordination is generally not easily made. The reason for this is ascribed to
the Jahn-Teller effect arising from the unequal occupation of the Eg pair of orbitals
(dz2 and dx2-y2) when a d9 ion is subjected to an octahedral crystal field. The d9
configuration makes CuII subject to Jahn-Teller distortion if placed in an environment
of cubic (i.e. regular octahedral or tetrahedral) symmetry, and this has a profound
effect on all its stereochemistry. The d9 configuration can be thought of as an
inversion of d1, relatively simple spectra might be expected, and it is indeed true that
the great majority of Cu (II) compounds are blue or green because of a single broad
absorption band in the region 11000-16000 cm1-, however as already noted, the d9ion
in characterized by large distortions from octahedral symmetry and the band is
unsymmetrical, being the result of a number of transitions which are by no means
easy to assign unambiguously. The free-ion ground 2D term is expected to split in a
crystal field in the same way as the 5D term of the d4ion, (2Eg→2T2g).and a similar
interpretation of spectra is like wise expected. Unfortunately this is now more difficult
because of the greater overlapping of bands which occurs in the case of Cu (II). The T
ground term of the tetrahedrally coordinated ion implies on orbital contribution to the
magnetic moment, but the E ground term of the octahedrally coordinated ion because
of ‘mixing’ of the excited T term into the E ground term with compounds whose
Chapter One Introduction
37
geometry approaches octahedral having moments at the lower end, and those with
geometries approaching tetrahedral having moments at the higher end, but their
measurements can not be used diagnostically with safety unless supported by other
evidence.(64,65)
Copper (I1) chloro complex formation was examined in basic melt by using
absorption spectroscopy. Potentiometric titration experiments involving the Cu (II)/Cu
(I) couple were attempted, but the resulting titration data were irreproducible and
could not be analyzed in a meaningful way. The exact cause of this behavior could not
be determined. Figure 16 shows absorption spectra for copper (I1) dissolved in 47.3
mol % AlC13,-MEIC melt and DMF containing either LiCl or MEIC as the chloride
source. The spectrum for copper (I1) in DMF, containing high concentrations of LiCl,
exhibits maxima at 295 and 411 nm. These absorption maxima have been reported for
Copper (I1) in propylene carbonate and dimethyl sulfoxide, also. They are
characteristic of the unsolvated CuC142- ion. A spectrum of copper (I1) in 47.3 mol %
AlC13,-MEIC exhibits maxima at 293 and 407 nm. The molar absorptivity for the
407-nm peak compares favorably with that for the 410-411nm peaks found in DMF
and other solvents the peak at 293 nm, observed for copper(I1) in basic melt, seems
distorted and exhibits a significantly higher absorbtivity than that observed for the
294-295-nm peaks in organic solvents containing LiCl. An absorption spectrum for
Copper (I1) in DMF with MEIC as the source of chloride ion exhibits maxima at 282
and 408 nm and an ill-defined shoulder near 292 nm as in (Figure 1-15).(72)
Anhydrous Copper (II) sulfate dissolved readily in the acetamide-potassium
nitrate eutectic to give blue-green solutions. Spectroscopy showed two bands at ν =
20700 cm−1 (ε= 43 mol−1 dm3 cm−1) and ν = 14000 cm−1 (ε = 73mol−1 dm3 cm−1)
with a shoulder at ν = 11000 cm−1. The copper cations appear to be in an
octahedral coordination, though with a higher extinction coefficient, and thus
probably a lower symmetry (perhaps D4h), than those found for the regular
octahedral complexes in aqueous solution (band maxima at ν = 12600 cm−1, ε = 12
mol−1 dm3 cm−1), or in anhydrous alkali metal nitrate melts (ν = 13000 cm−1, ε = 14
mol−1 dm3 cm−1), because in this binary eutectic coordination is by both acetamide
Chapte
and nitr
due to sm
Fig (1-15):melt, 6.8 ×5.0 M in C
1.13. Th
melts.
The
(LiCl/K
oxidatio
octahedr
(III) in
ν1=1850
same m
coordina
tetrahed
370-450
indicatin
range o
ν3=1470
LiCl/KC
er One
rate. The h
mall amou
:- Copper (I1×105- M in CoCopper (I1); -
he coord
coordinat
KCl) has b
on state (
ral coordi
chloride
00 cm-1, ν2
melt at 40
ation and
dral coord
0oC show
ng probab
of 447oC t
00 cm-1,
Cl melt sh
higher ene
unts of a s
1) absorptionopper (I1); ---, DMF + 0.4 M
dination c
tion chem
been the s
II) and (I
ination by
melt at
2=12500 c
00oC show
at 1000oC
ination, w
wed one a
bly tetrahe
to 480oC
also indic
howed the
ergy band
square pyr
n spectra in so-, DMF + 2.7 M MEIC, no
chemistr
mistry of
ubject of
III) cation
y distortion
400oC sh
cm-1 indic
wed one
C showed
while Fe (
absorption
dral coord
showed
cating tet
an absorp
38
d is anoma
ramidal (C
olvents contaM LiC1, 5.4
o Copper (I1).
ry of tran
transition
many stu
ns of the
n and form
howed tw
cating an o
band at
one band
III) in the
n band at
dination. C
three ba
trahedral
ption of th
alous for a
C4v) compl
aining chlorid ×105- M inC.(72)
nsition m
n metal
udies whic
first row
mation of
wo bands
octahedral
ν=5100 c
at ν=6000
e same m
t 25000
Co(II) in c
ands ν1=16
coordinat
ree bands
a D4h com
lex.(67)
de ion: —, 47Copper(I1); -.
metal cati
cations in
ch showed
w transitio
f tetrahedr
as maxim
l coordina
cm-1 indi
0 cm -1 in
melt at tem
cm-1 in a
chloride m
6700 cm-1
tion. In a
at ν=1530
Introd
mplex and
7.3 mol % Al-, DMF + 3.0
ions in c
n chlorid
d the devi
on elemen
ral geome
ma absor
ation, Fe(I
cating tet
ndicating d
mperature r
a shoulde
melt at tem1, ν2=1515
addition N
00 cm-1, ν
duction
d may be
lC13-MEIC 0 M MEIC,
chloride
de melts
iation of
nts from
etries. Cr
rption at
II) in the
trahedral
distorted
range of
er shape
mperature
50 cm-1,
Ni(II) in
ν= 14200
Chapter One Introduction
39
cm-1, ν=8000 cm-1 at 700-1000oC, indicating distorted tetrahedral coordination, and
in CsCl melt at 864oC also showed three absorption bands at ν=15700 cm-1,
ν=14000 cm-1, ν=7500 cm-1 indicating similar tetrahedral coordination. Cu (II) in
the chloride melt (LiCl/KCl) at 400oC showed one absorption band at ν=9500 cm-1
indicating tetrahedral coordination.(1)
In the case of Ni(II) the three bands were initially assigned to a distorted
tetrahedral coordination though in the light of more detailed measurements over a
temperature range, the coordination ratio tetrahedral/octahedral centers was about
to increases with potassium chloride concentration (in LiCl/KCl mixture) as well
as with increasing temperature.(1)
1.14. Determination of spectra.
1.14.1. Orgel diagrams.
Orgel diagrams are useful for showing the number of spin-allowed absorption
bands expected, and their symmetry state designations, for Td and weak field Oh -
complexes (d5 not included). The two diagrams (one for d1, 4, 6, 9 the other for d2, 3, 7,
8) pack a lot of information in very little space. Nevertheless, there are three major
limitations to using Orgel diagrams:
(1) They are restricted to weak field/high spin conditions and offer no
information for d4, 5, 6, 7 strong fields/ low spin cases.
(2) They only show symmetry states of same highest spin multiplicity
(3) They are qualitative, and therefore energy values cannot be obtained or
calculated from them.(74)
Chapte
Fig. (
Fig. (1-17
1.14.2.
In con
from a
from Or
horizont
and dire
d 7, case
these di
er One
(1-16):- The
7):- The sp
Tanabe-S
ntract to O
configurat
rgel diagra
tal base, w
ect ground
es for both
iagrams a
e splitting o
plitting of fr
Sugano d
Orgel diag
tion, not j
ams becau
whereas O
d symmetr
h low and
appear to
of free-ion
ree-ion F an
diagrams
grams Tan
just terms
use they u
Orgel diagr
ry states b
d high spin
be disco
40
D terms in
nd P term i
s.
nabe-Sugan
s of highe
use the gro
rams place
elow it. In
n symmetr
ontinuous
n octahedra
in octahedr
no diagram
est spin-mu
ound symm
e the paren
n addition
ry states a
- having
al and tetra
ral and tetr
ms contain
ultiplicity
metry state
nt term in
n the splitt
are includ
two part
Introd
ahedral fiel
rahedral fi
n all term
y. They als
e as a stra
n a central
ing of d 4,
ded. Conse
ts separat
duction
ds. (74)
eld.(74)
ms arising
so differ
aight-line
location
, d 5, d 6,
equently,
ted by a
Chapter One Introduction
41
vertical line. The left part pertains to the weak field /high spin condition and the
right to strong field/ low spin condition.
First note why d 1 and d 9 cases have no T-S diagrams. A term description for an
atom/ion is more informative than its electron configuration because terms account
for e-e repulsion energies. However there is no e-e repulsion for one "d" electron
so the d 1 configuration gives rise to a single term, 2D. In Oh and Td ligand fields
this single term is split into T2g, Eg, or E, T2 symmetry states. Only one absorption
band is expected and energy of the observed band gives the Δo or ΔTd value
directly. No calculations are necessary, so no T-S diagram for d 1 (and d 9).(74)
Chapter One Introduction
42
1.14. Aim of the present work.
New class of room temperature ionic liquids (RTILs) based on choline chloride
have been prepared.(59) These ionic liquids offered promising ease of handling and
wide field of application as solvents and reacting media.(61) The behavior of some
transition metal chlorides were investigated by ultraviolet visible spectroscopic
measurements in two types of RTILs melts. The cationic metal are Cr3+, Fe3+, Co2+,
Ni2+ and Cu2+ while the RTILs are choline chloride/ urea and choline chloride/
malonic acid mixtures which have freezing point of 12oC and 10oC respectively.
The solutions of these transition metals in ionic liquids will be studied by
ultraviolet visible spectroscopy, infrared spectroscopy, conductivity and magnetic
susceptibility, to determine their behavior and coordination in solution.
Chapter Two Experimental Part
43
Chapter Two
Experimental Part
2.1. Chemicals and Instruments.
2.1.1. Chemicals.
All the chemicals used in this work were of highest purity available and
without further purification. Table (2-1) shows the reagents, their purities and
suppliers.
Table (2-1):- The chemicals used in this work with their purities and suppliers.
Chemical compounds Purity % Company
Choline chloride 99 Merck
Urea 99.5 Merck
Malonic acid 99 Chem-Supply
Ammonium nitrate 99.9 Baker Analyzed Reagents
Chromium(III) chloride
hexahydrate
97 Chem-Supply
Iron(III) chloride tetrahydrate 99 BDH
Cobalt(II) chloride
hexahydrate
98 Chem-Supply
Nickel(II) chloride
hexahydrate
99 BDH
Copper(II) chloride dihydrate 95 Fluka
Cobalt(II) nitrate hexahydrate 98.5 BDH
Chapter Two Experimental Part
44
2.1.2. Instruments.
2.1.2.1. Electronic absorption spectra measurements.
The electronic spectra of complexes were obtained by using SHIMADZU
(Japan), UV-Vis 1650 PC Ultraviolet Visible Spectrophotometer at room
temperature using quartz cells of 1.0 cm length and ionic liquids as solvents in the
range of wave length 190-1100 nm. The preparations of samples are described in
section (2.5).
2.2. Setting of heating apparatus.
The heating apparatus that was used in the preparation of the ionic liquid is
consisted of: -
1- Hot plate.
2- Stand.
3- Clamp.
4- Thermometer.
5- Oil bath.
6- Reaction beaker.
7- Glass rod.
2.3. Drying method.
The drying method was carried out by placing appropriate weight of samples in
the oven under temperature range of 70 -75oC; the dried compound was kept in
sealed the desicator for further use.
Chapter Two Experimental Part
45
2.4. Preparation of room temperature ionic liquid.
2.4.1. Preparation of Choline chloride/Urea room temperature ionic
liquid.
Solid mixture of dry choline chloride and urea in mole ratio of 1:2 was milled,
and well mixed and heated gradually to 75oC for a period of 20 min. using the
apparatus described in (section 2.2) with continuous gentle stirring until both
choline chloride and urea melted together producing colourless liquid, the mixture
was gradually cooled down to room temperature, it was more viscose than water.
Finally the prepared melt was kept in sealed dry desiccator for further use. (58, 61)
2.4.2. Preparation of Choline chloride/Malonic acid room temperature
ionic liquid.
Solid mixture of dry choline chloride and malonic acid in mole ratio of 1:1 was
milled, and well mixed and heated gradually to 85oC for a period of 20 min. (62)
using the apparatus described in (section 2.2) in the same way in (section 2.4.1).
2.5. Preparation of metal chloride ionic liquid mixtures.
The fallowing general procedure was used for the preparation of the solution of
metal chlorides salts to ionic liquids:-
A known weight of dry sold compound [Table (2-2)] was added to appropriate
volume of ionic liquid at room temperature, the mixture was left to stand for 72
hours with occasional stirring to enhance dissolution of materials. The changes
during this period were recorded such as colour changes and viscosities, however,
no precipitation or gas evolution were noticed through out the entire experiments
of this work.
In addition, during the investigation of coordination reaction, ammonium nitrate
was added to above solution and any noticeable changes were recorded.
Chapter Two Experimental Part
46
All prepared solutions of metal chloride in ionic liquids of chlorine chloride/ urea
with and without the addition of ammonium nitrate and choline chloride/ malonic
acid were investigated by Ultraviolet Visible Spectroscopy and the obtained
spectra were analyzed. Table (2-2):- Concentration of metal chloride in aqueous solution and room temperature ionic liquid of choline chloride/urea and choline chloride/malonic acid.
Metal chloride
Weight (g.)
Concentration in
aqueous solution
(M)
Concentration in
choline chloride/urea
ionic liquid (m)
Concentration in
choline
chloride/malonic acid
ionic liquid (m)
CrCl3.6H2O 0.3 g. 0.1 M 7.4×10-2 m 7.5×10-2 m
FeCl3.4H2O 0.2 g. 0.1 M 6.6×10-2 m 6.9×10-2 m
CoCl2.6H2O 0.25 g. 0.1 M 6.9×10-2 m 7.4×10-2 m
NiCl2.6H2O 0.4 g. 0.1 M 1.1×10-1 m 1.2×10-1 m
CuCl2.2H2O 0.2 g. 0.1 M 7.7×10-2 m 8.03×10-2 m
M: - molarty, which is number of moles of solute per volume of solvent in L unit.
m: - molalty, which is number of moles of solute per weight of solvent in Kg. unit.
2.6. Equations.
The equations that have been used in the determination of spectra of each
complex were different from metal to another according to the splitting of terms
for dn configurations as shown in Figure (2-1) and in Table (2-3).
Fig. (2-1):- Splitting of terms for dn configurations in to (a) and (b).(65)
Chapter Two Experimental Part
47
Table (2-3):- Splitting of terms elected dn configuration.(65)
Configuration of free-ion
Ground state of free-ion
Energy level diagram
Predicted in Figure (23)
d3 4F Inverted (b) ν1=10 Dq ν2=18Dq-x ν3=12Dq+15B`+x
d5 6S No splitting d7 4F (b) ν1=8Dq
ν2=18Dq ν3=6Dq+15B`
d8 3F Inverted (b) ν1=10 Dq ν2=18Dq-x ν3=12Dq+15B`+x
d9 2D Inverted (a) ν1=10Dq
Also Jorgenson rule have been used to measure the tendency of metal ion to
form a complex, the (g) factors provide an estimate of the value of 10Dq for an
octahedral complex when combined with the (f) value for the appropriate
ligands:(85)
10Dq = f ligand × g ion ×10 3 _____________________________ (2-1)
Table (2-4):- Values of g and f factors for metal ions and various ligands.
Transition metal ion Value of (g) factor Various ligands Value of(f) factor
Cr(III) 17.4 Cl- 0.78
Fe(III) 14.0 CO(NH2)2 0.92
Co(II) 9.00 H2O 1.00
Ni(II) 8.7 (CH2COOH)2 0.96
Cu(II) 12.5 NO3- 0.88
In units of k K (=1000 cm-1).
If all three transitions are observed, it is a simple matter to assign a value to B`,
since the following equation must hold; (B` is in cm-1 units).
Chapter Two Experimental Part
48
15B`= ν3+ ν2–3 ν1 __________________________________________ (2-2)
The nephelauxetic ratio β is given by:
β= B`/B. _________________________________________________ (2-3)
Where ‘B’ is Racah parameter.
And by using equation (2-4) to find the ΔT,
10Dq Td = 4/9 (10Dq Oh) _________________________________________ (2-4)
2.7. Example of calculations.
Consider an example application of a Tanabe-Sugano diagram for a d3 case. The
absorption spectrum of a Cr 3+ solution is shown below in Table (2-5). For example
the mathematical determination of spectra of CrCl3.6H2O in choline chloride/urea
ionic liquid at room temperature was taken. Table (2-5):- Peaks positions of Chromium (III) ions in room temperature ionic liquid of Choline chloride/Urea solution.
Ion ν1 cm-1 ν2 cm-1 ν3 cm-1 ν2/ ν1 ν 1/ ν2 Δ/B
Cr+3 16447 22727 ? 1.38 0.724 26.00
Two bands are observed within the range of measurement. They have a maxima
at about 16447 and 22727 cm-1. These are spin-allowed LaPorte-forbidden d-d
transitions. Chromium is in the 3+ oxidation state, so this is a d3 system. Reference
to an Orgel diagram informs that three d-d bands are expected and they can be
assigned as:
ν1= 4A2g (F) → 4T2g (F)
ν2= 4A2g (F) → 4T1g (F)
ν3= 4A2g (F) → 4T1g (P) Δo is taken to equal the absorption energy of 16447 cm-1, and the intersect on the
x-axis of Tanabe-Sugano diagram equal Δo/B and by drawing a vertical line from
Chapter Two Experimental Part
49
this point will intersect with other allowed electronic states of Tanabe-Sugano
diagram for d3.
There are several “goals” sought in analyzing spectra using Tanabe-Sugano
diagrams: (99)
(1) To make correct band assignments. The two bands observed could be the first
and second, or the second and third. Their assignment cannot be made by
inspection.
(2) To determine the magnitude of the Ligand Field splitting parameter, ∆o.
(3) To determine the magnitude of the e-e repulsion parameter (called a Racah B
parameter).
Assumes bands are the first and second (so third band is not observed). Compare
the two results. Tanabe-Sugano diagrams are unit less graphs showing energy
ratios. The abscissa shows values for the ratio ∆o/B (i.e., Ligand Field splitting
parameter / e-e repulsion parameter) and the ordinate shows values for the ratio
(Eν/ B) (i.e., absorption band energy / e-e repulsion parameter).(74,75)
Step (1): The calculated ratio of experimental band energies is: E (ν2) / E (ν1), find
chart lines for first and second symmetry states and record their values at each
abscissa mark.
None of these chart ratios come close to the experimental ratio of (1.38).
Step (2): Find the specific value for ∆o/B, having a chart value for ν2/ν1 that
matches the experimental ratio of 1.38.
When ∆o/ B = 30, chart ν2/ν1 = 1.40, and at ∆o/ B = 20, chart ν2 /ν1 = 1.30.
Verify (by extrapolation) that at a value of ∆o/ B = 26, the chart ratio for ν2/ν1=
1.38.
Step (3): Find the intersection of chart lines for 4T2g and 4T1g at ∆o / B = 26,
And record their values on the ordinate E (νn)/ B.
For 4T 2 g at ∆o / B = 26, then E (ν1) / B = 26, and
Chapter Two Experimental Part
50
For 4T1g at ∆o / B = 26, then E (ν2) / B = 37
Step (4): Use experimental values of E (ν1) and E (ν2) to obtain values for Racah B:
B [via E (ν1)] = 16447 cm-1 / 26 = 623.6 cm-1
B [via E (ν2)] = 22727 cm-1 / 37 = 614.3 cm-1
Then, B avr. = 619 cm-1.
The two determinations are in acceptable agreement. Furthermore, in free Cr
(III) ion, Racah B = 918 cm-1, a larger value. It is anticipated that e-e repulsion
should be reduced in compounds because metal electron density becomes
delocalized (as a result of bond formation to ligands) and metal e-e repulsions are
decreased correspondingly.
Step (5): Use this value for Racah B to determine ∆o, the Ligand Field splitting
parameter.
If ∆o / B = 26, and B = 619 cm-1, then ∆o = 16447cm-1.
All calculated values are reasonable. But what about the missing third band?
Step (6): Three bands are expected, but only two are observed. Fitting the first two
absorption bands to the Tanabe-Sugano chart, according to the assignments
produced good results. Information gained can be used to calculate the energy
expected for the missing third transition:- ν3= 4A2g (F) → 4T1g (P)
At ∆o / B = 26, the chart value for E(ν3) / B = 56
Using the value already obtained for B produces the result E (ν3) calc’d = 56.00 ×
618 cm-1= 34608 cm-1.
Is this reasonable?
(1) Refer to the diagram for the recorded spectrum and note that it only extends to
35,000 cm-1. This is the upper energy limit of the instrument used to measure the
spectrum.
(2) Notice that a small error in reading spectral band energies changes the
experimental ν2/ν1 ratio which in-turn changes the matching ∆o/B ratio on the
Tanabe-Sugano diagram.
Chapter Two Experimental Part
51
This means a small change in ∆o / B ratio will cause a large change in Eν / B for
ν3. Either or both of these factors could be responsible for third band not being
observed. Observed at 16447 cm-1, 22727 cm-1 (third band not observed) Ligand
Field parameters: ∆o = 16447 cm-1. Racah (B) = 619 cm-1 (free ion B = 918 cm-1).
And if the three expected bands were observed then by using the equations of
splitting of term for d3 in Table (2-3), we can determine the B` and β as following:
ν1=10 Dq.
16447 cm-1= 10 Dq
Dq = 1644.7 cm-1.
ν2=18Dq-x
22727 cm-1 = 18(1644.7 cm-1) – x
x = 6877.6 cm-1.
ν3=12Dq+15B`+x
34608 cm-1 = 12 (1644.7 cm-1) +15B`+6877.6 cm-1
B`= 533 cm-1
Or we can use the following equation to determine the B`
15B`= ν3+ ν2−3ν1
15B`=34608 cm-1+22727 cm-1-3(16447 cm-1)
B`=533 cm-1
To determine the nephelauxetic ratio β:
If β= B`/B, Then β=533 cm-1/619 cm-1
And β=0.68.
Chapter Three Results and Discussion
52
Chapter Three
Results and Discussion
3.1. Introduction.
This project relates to ionic compounds and methods for their preparation. In
particular the project relates to ionic compounds which are liquid at relatively low
temperatures, generally below about 100°C, and preferably below about 60°C and
more preferably which are liquid at or near to ambient temperature. In recent years,
ionic compounds have been made which are liquid at relatively low temperatures,
for example, trimethyl phenyl ammonium chloride with aluminium trichloride. The
resulting ionic compound has a low freezing point (around -75°C), but suffers from
the same water sensitivity, because of the presence of aluminium trichloride. There
is much current interest in the field of ionic liquids. Such systems, which are
examples of molten salts, have a number of interesting and useful chemical
properties, and have utility, for example, as highly polar solvents for use in
preparative chemistry, and as catalysts. They also have particular application in
electrochemistry, for example in batteries, fuel cells, photovoltaic devices and
electrodeposition processes, for example in baths for the electroplating of metals.
Ionic liquids have very low vapour pressure and thus, in contrast to many
conventional solvents, are very advantageous in that they produce virtually no
hazardous vapours. They are therefore advantageous from a health, safety and
environmental point of view.(63)
A new class of room temperature ionic liquid based on choline chloride have
recently been prepared which have an advantageous over previous room
temperature ionic liquid. The major advantageous are water insensitive which
made them attractive RTIL for chemical and industrial applications. The potential
of their advanced properties also arise from the availability, reasonable price and
green characterization. The emerged RTILs have caught the attention of some
Chapter Three Results and Discussion
53
researchers to establish their chemical, physical and industrial scope of
understanding.
Therefore, it is not surprising, in addition to its own right, to study the
characterization of some transition metal ions (Cr3+, Fe3+, Co2+, Ni2+ and Cu2+) to
determine their behaviour, particularly, coordination, interaction in such new clan
of solvents.
3.2. Electronic absorption of transition metal chloride in room
temperature ionic liquids. Electronic absorption spectra of transition metal complexes are usually
attributed to the partially filled d-orbital of the metal. The energy required for such
transitions is that of the near ultraviolet and visible region. Charge transfer spectra
are due to transitions between metal and ligand. Studies of electronic spectra of
complexes help in the determination of structure of the complexes through the
electronic interaction of the metal d-orbital and ligand orbital. (64, 65) In our work the
spectra were investigated in the range (190-1100) nm, using two types of ionic
liquids as solvent.
3.2.1.Spectroscopic investigation of Chromium (III) chloride solutions.
3.2.1.1. Chromium (III) chloride in choline chloride/urea room
temperature ionic liquids.
Chromium (III) chloride hexahydrate was found to be completely soluble in
choline chloride/ urea room temperature ionic liquid after 72 hours in
concentration of 7.4×10-2m at room temperature see Table (3-1), producing a deep
green colored solution. The viscosity of the choline chloride/urea ionic liquid
before and after the dissolution of the salt seemed visually to have close viscosity.
The recorded ultraviolet visible spectrum of the solution showed two absorption
bands at 16447 cm-1(608 nm), at 22727 cm-1 (440 nm) and a shoulder at 14502cm-1
(689.56 nm). The U.V.-Vis. spectrum Figure (3-1) indicates that the Chromium
(III) ion is consistent with octahedral coordination, as shown in Figure (3-1).
Chap
Fig. (3-1)temperatu The t
region
electron
ground
may be
[Figure(
And the
forbidde
Fig
pter Thre
:- The U.V.-ure ionic liq
third band
was dedu
nic configu
state of C
assigned
( 1-17)] as
e appeared
en transitio
g. (3-2):- Cr
ee
-Vis spectruquid.
d which n
uced by
uration an
Cr (III) ion
d as illustr
s fallow :-
d shoulder
on from th
rystal field
um of Chrom
normally
using Ta
nd found
n 4F can
rated in F
ν1= 4A
ν2= 4A
ν3= 4A
on the fir
he ground
splitting of
54
mium (III) o
occurred
anabe-Sug
to be 34
split in oc
Figure (3-
A2g (F) → 4T
A2g (F) → 4T
A2g (F) → 4T
rst absorpt
d state of 4A
f 4F term of
Re
of (7.4×10-2m
at the hi
gano diag
733 cm-1.
ctahedral
-2) and ac
T2g (F) =10
T1g (F)
T1g (P)
tion peak w
A4g to the
f d3 ion oct
esults an
m) in Cholin
igh charg
gram Figu
. The term
crystal fie
ccording
0Dq
was assign
e doublet s
ahedral co
nd Discu
ne chloride/U
ge transfer
ure (3-6)
m symbol
eld and th
to Orgel
ned to be o
state 2G.
nfiguration
ussion
Urea room
r energy
) for d3
l for the
he bands
diagram
of a spin
n.(64)
Chap
It wa
for Chro
(III) ion
of as 12
be octah
surprisin
of urea
Howeve
nitrate
eutectic
octahedr
in the in
pter Thre
Fig. (s found th
omium (II
n in this eu
2500 cm-1
hedral coo
ng that the
than chl
er, the ba
(16760,22
s as in
raly coord
ntroduction
ee
(3-3):- Tanahat the ob
II) ion in
utectic we
and 1850
ordinated
e higher en
loride and
and energi
2450 cm-
Table (3-
dinated wi
n (section
abe-Suganoserved ba
other ioni
ere higher
0 cm-1, wh
with chlo
nergy obs
d it is in
ies were 1) or but
-1), wher
ith six ace
1.12.1.).(6
55
o diagram and energi
ic liquids.
than in pu
here the C
oride ions
served in t
agreeme
close to t
tyramide-s
re Chrom
etamide o67)
Re
for d3 ion ces were c (67) The b
ure chlorid
Chromium
s.(1)[Table
this work e
ent with t
those foun
sodium n
mium (III)
or six nitra
esults an
configuratilose to th
band energ
de melt lik
m (III) ions
(3-1)] If
explains t
the spectr
nd in ace
nitrate (16
) ions a
ate ligands
nd Discu
ion.(73) ose band
gies of Ch
ke LiCl/K
s were ass
so, then
the stronge
rochemica
etamide-po
6600,2250
assigned
s, as state
ussion
energies
hromium
KCl melt,
signed to
it is not
er nature
al series.
otassium
00 cm-1)
to have
ed earlier
Chap
Fig. (3-4)choline ch
The b
this wo
aqueous
stronger
the abso
liquid w
(3-2). F
behavio
eutectic
Calcu
Chromiu
urea mo
the latte
[I- < Br
CH2(CO
≈ en (et
NO2- < P
pter Thre
):- The U.Vhloride/ urea
band energ
ork were
s solution
r than ure
orbance an
were found
Figure (3-4
or of Chro
.
ulations u
um (III) io
olecules. T
er ion is at
r- <S2- <
OO)22- ≈ C
thylenedia
PPh3 < CN
ee
V.-Vis. Specta eutectic so
gies of ch
found to
n Table (3
ea molecu
nd the mo
d to be sm
4) and Ta
omium (II
using Jor
ons in this
This is exp
the lowes
SCN- ~
C2O42- < H
amine) <
N- ≈ CO].
tra of CrClolution (7.4×
hromium (
be shifte
3-1). This
ule accord
lar absorp
maller than
able (3-1)
I) ion in
genson R
s eutectic
pected in v
st weak en
Cl- <NO3
2O < NCS
bipy (2,2
.
56
l3.6H2O (0.1×10-2m), (cu
(III) ions
ed to low
s might b
ing to the
ptivity of C
n that obs
which sh
aqueous s
Rule of a
were octa
view of th
nd of the s
3- < N3
- <
S- < CH3C
'-bipyridin
Re
1 M) in aquurve B) at ro
in choline
wer energi
be due to
e spectroc
Cr (III) in
served in
how the d
solution an
average en
ahedral co
he stronge
pectroche
< F- < CO
CN < EDT
ne) < ph
esults an
ueous solutioom temper
e chloride
ies than t
o water m
chemical s
n choline c
aqueous s
difference
nd in cho
nvironmen
oordinated
er urea tha
emical seri
O(NH2)2 <
TA4- < NH
hen (1,10-p
nd Discu
on (curve Arature.
e/urea obs
those obta
molecule w
series. In
chloride/ur
solution se
in the ill
oline chlor
nt sugges
d with 6 lig
an chloride
ies.
< OH- ~
H3 ≈ pyr(p
phenanthr
ussion
A), and in
served in
ained in
which is
addition
rea ionic
ee Table
lustrated
ride/urea
sted that
gands of
e ions as
ONO- <
pyridine)
roline) <
Chapter Three Results and Discussion
57
The consistency of this interpretation can be checked and a value was obtained
for the ‘interelectronic repulsion parameter’ B, by fitting the ratio of ν1/ν2 from the
appropriate Tanabe-Sugano diagram, Figure (3-3). The obtained value of B equals
to 632 cm-1 as in [Table (3-1)], which is smaller than that for free ion 918cm-1 or
that of the metallic ion in the aqueous solution 918cm-1. The low value of B
normally arises when the ligands have more expansion space for d-electron as in
urea. And indicates coordination via covalent rather than ionic bonds where high
value ought to be formed. The solution of Chromium (III) chloride hexahydrate in
choline chloride/urea room temperature ionic liquid did not showed any observed
changes in ultraviolet visible spectra when the solution was stored for 40 days,
indicating that the solution already reached equilibrium and was stable over this
period of storing time.
3.2.1.2. Substitution Reaction with Ammonium Nitrate.
When ammonium nitrate (NH4NO3) was added to the chromium (III) chloride
hexahydrate in choline chloride ionic liquid solution at room temperature, it
showed no changes either in color or in the apparent viscosity of the solution.
However, when the mixture was examined by ultraviolet visible spectrum of the
Chromium (III) ion, the observed peaks shifted to higher energies than those
obtained in pure melt and have the following values 16750 cm-1 (597 nm) and
23529 cm-1(425 nm) and a shoulder at 14502 cm-1 (689.56 nm). The addition of
nitrate also showed an increase in the maximum intensity of those bands in
comparison with those obtained in choline chloride/urea alone as shown in Figure
(3-5). The changes might indicate some interaction of the added nitrate with
Chromium (III) ions. The third energy band was calculated from the Tanabe-
Sugano diagram for d3 ion configuration Figure (3-3) and it was found to be equal
to 36551 cm-1.
Calculation using Jorgenson Rule of average environment was found to be in
agreement with replacement of two urea molecules by two nitrates in octahedral
Chap
coordina
the spec
indicate
molecul
with in
ionic the
symmet
Fig. (3-5)2m) (curvroom tem
The C
ionic liq
indicatin
The R
equal to
even the
that val
liquid a
interacti
pter Thre
ation; nitr
ctrochemi
es some di
les with tw
the plane
e preferab
trical comp
:- The U.V.ve A), and i
mperature.
Chromium
quid did n
ng a stable
Racah par
o 664.56 c
e hydrated
lue calcula
alone 632
ion with th
ee
rate is con
cal series
istortion in
wo nitrates
e of the ge
ble one is t
plex, as sh
.-Vis. Spectrin choline c
m (III) chlo
not show
e complex
rameter B
cm-1, this
d chromiu
ated for c
2 cm-1, th
he added n
nsidered to
, as show
n symmet
s see Tabl
eometry, a
the replace
hown in F
ra of dissolvchloride/ure
oride solut
any chang
x over this
B was calc
value is m
um (III) co
chromium
he higher
nitrates an
58
o be stron
wed in Fig
try due to
e (3-2). Th
as the for
ement tha
igure (3-5
ved CrCl3.6Hea with add
tion in cho
ges in ultr
s period.
culated fo
much sma
omplex 71
m (III) com
r value in
nd the form
Re
nger than u
gure (3-5)
the substi
his distort
rmed bond
t occurred
5).
H2O ion in ition of NH
oline chlo
raviolet v
or the com
ller than t
13 cm-1. H
mplex in
ndicate th
med bond
esults an
urea as lig
. The inc
itution rea
tion might
ds with th
d on the z-
choline chloH4NO3 salt (
oride/urea
isible spe
mplex and
that of fre
However, i
choline ch
hat there
s were ion
nd Discu
gand acco
rease abs
action of t
t either be
he nitrates
-axis to for
oride/urea o(curve B) eu
room tem
ctra after
d was foun
ee ion 918
it was hig
hloride/ur
is some
nic which
ussion
ording to
orbtivity
two urea
z-out or
is more
rm more
of (7.4×10-
utectics at
mperature
40 days
nd to be
8 cm-1 or
gher than
rea ionic
kind of
increase
Chapter Three Results and Discussion
59
the value of B, the higher value normally indicated a more repulsion in d-electron,
due to more expansion of d-electron over a large cloud space.
3.2.1.3. Chromium (III) chloride in Choline chloride/Malonic acid room
temperature ionic liquid solution.
Chromium (III) chloride hexahydrate was found to be completely soluble in
choline chloride/ malonic acid ionic liquid at room temperature after standing time
of 72 hours in concentration of 7.5×10-2m, producing a greenish grey solution
without apparent chang in viscosity of the choline chloride/ malonic acid ionic
liquid. The ultraviolet visible spectrum of the solution showed two absorption
bands at 15576 cm-1 (642 nm), 21141 cm-1 (473 nm) and a shoulder at 14372 cm-
1(696 nm) Figure (3-6). The energies of these bands agree with those of octahedral
coordination of Chromium (III) complexes.
The third spectra band of Chromium (III) complex which is expected to have
higher energy than ν1 and ν2 was calculated with the used of Tanabe-Sugano
diagram of d3 ion configuration Figure (3-3) and found to equal to 33994 cm-1 (294
nm) at the ultraviolet region. The bands were suggested to be due to the fallowing
transition:-
ν1= 4A2g (F) → 4T2g (F) =10Dq
ν2= 4A2g (F) → 4T1g (F)
ν3= 4A2g (F) → 4T1g (P)
While the shoulder that appeared was assigned to be of a forbidden transition from 4A4g ground state to 2G state.
Chap
Fig.(3-6):acid room Calcu
choline
of the s
coordina
ligand to
most su
chloride
made th
The
chloride
chloride
chloride
Table (3
chloride
pter Thre
- The U.V.-m temperatu
ulations u
chloride/m
ix chlorid
ated to Cr
o the Cr (I
uitable rep
es are on t
he formed
band ene
e/urea eute
e/ urea ion
e/urea eute
3-2), that
e/urea and
ee
-Vis. Spectrure ionic liqu
using Jorg
malonic a
de ions in
r (III) ion
III) octahe
placement
he z-axis
complex m
ergies are
ectic, this
nic liquid
ectic whic
shows th
d choline c
um of CrCluid.
genson Ru
cid ionic
the octah
as a biden
edral struc
t of chlori
and the ot
more relax
shifted
indicate
. The abs
ch indicate
he deferen
chloride/m
60
l3.6H2O ion
ule of aver
liquid ind
hedral coor
ntate ligan
cture migh
ide ions m
ther two a
xed. As su
to lower
that malo
orbance w
es less dis
nce betwee
malonic aci
Re
of (7.5×10-2
rage envir
dicate that
rdination.
nd. The wa
ht occurred
might be
are on the
uggested in
energy a
nic acid h
was found
storted, as
en the spe
id ionic liq
esults an
2m) in Choli
ronment fo
malonic a
Therefor
ay of attac
d in differ
occurred
plane of x
n Figure (
as compa
has weak
d to be les
s shown in
ectra of C
quids.
nd Discu
ine chloride
for Cr (III)
acid repla
re, maloni
ched malo
rent direct
when tw
x and y ax
(3-8).
ared with
field than
ss than in
n Figure (
Cr (III) in
ussion
e/ Malonic
) ions in
aced four
c acid is
onic acid
tions, the
o of the
xis which
choline
n choline
n choline
3-7) and
n choline
Chap
Fig. (3-7)A), and in
The
obtained
fitting th
found to
cm-1) an
cm-1), b
the valu
The a
This ph
delocali
only the
expandi
in the co
The c
sequenc
[Free
urea > c
pter Thre
:- The U.V.n choline chl
consisten
d for the
he ratio of
o be equal
nd as well
ecause of
ue of B.
apparent va
henomeno
ization of
e metal bu
ing” the av
omplex) is
calculated
ce as ligan
ion > H2
choline chl
ee
-Vis. Spectrloride/malon
ncy of thi
interelect
f ν1/ν2 the
l to 539.6
l as less t
f the more
alue of B
on is kn
the meta
ut the liga
verage int
s smaller.
values f
nd changes
2O > choli
loride/ ma
ra of CrCl3.nic acid of (
is interpr
tronic rep
appropria
cm-1, the
than that C
expansion
in comple
nown as
l electron
and as we
terelectron(65)
for Racah
s:
ine chlori
alonic acid
61
6H2O ion in(7.5×10-2m) (
etation ca
pulsion pa
ate Tanabe
value of B
Cr (III) in
n of d-ele
exes is alw
nephelau
ns over mo
ll. As a re
nic repulsi
h parame
de/ urea w
d].
Re
n choline ch(curve B) eu
an be ca
arameter o
e-Sugano
B was less
n choline
ctron char
ways smal
uxetic eff
olecular o
esult of th
ion is redu
ter could
with added
esults an
hloride/urea utectics at ro
lculated a
or Racah
diagram f
s than tha
chloride/u
rge on com
ller than th
fect and
orbitals th
his deloca
uced and B
d be put
d nitrate >
nd Discu
of (7.4×10-2
oom temper
and a va
paramete
figure (3-6
t for free
urea soluti
mplexation
hat of the
is attrib
hat encom
alization o
B (represe
in the fo
>choline c
ussion
2m) (curve rature.
alue was
er B, by
6). It was
ion (918
ion (632
n reduce
free ion.
buted to
mpass not
or “cloud
enting B
ollowing
chloride/
Chapter Three Results and Discussion
62
Table (3-1):- Electronic spectra for complexes of Chromium (III) ions in aqueous solution and various ionic liquids.
Solution of Chromium (III) in:
Absorption band cm-1
B (cm-1)
calculated 15B (cm-1)
calculated β
10Dq
Suggested structure
Reference
Aqueous Solution at room temperature
ν1=16474 ν2=23474 ν3cal.=37118
713.8 10707 1.043 16474
Oh (6 H2O) [Cr(H2O)6]3+
Present work
Choline chloride/Urea at room temperature
ν1=16447 ν2=22727 ν3cal.=34733
632 9480 0.873 16447
Oh (6 urea) [Cr(urea)6]3+
Present work
Choline chloride/Urea + NH4NO3
at room temperature
ν1=16750 ν2=23529 ν3cal.=36551
664.56 9967.5 0.986 16750
Oh (4 urea, 2 NO3
-) [Cr(urea)4(NO3)2]+
Present work
Choline chloride/ Malonic acid at room temperature
ν1=15576 ν2=21141 ν3cal.=33994
539.6 8094 1.04 15576
Oh (2 malonic acid, 2 Cl-) [Cr(malonic)2Cl2]-
Present work
LiCl//KCl at 400oC
ν1=12500 ν2=18500
-
-
-
-
Oh (6 Cl-) [CrCl6]3-
1
Acetamide-KNO3
at 90oC
ν1=16760 ν2=22450
-
-
-
-
Oh (6 acetamide)
67
Butyramide-NaNO3
at 140oC
ν1=16600 ν2=22500
-
-
-
-
Oh (6 butyramide)
67
Chap
From
ionic li
octahedr
with the
field tha
choline
where a
of that N
series, a
NO3- , b
molten
coordina
chloride
coordina
most sta
8).
Fig. (3-8)choline chionic liqu
pter Thre
Table (3-
iquid of
ral coordi
e spectroc
an chlorid
chloride/
a replacem
NO3- is in
as well as
but they w
salt prope
ation. And
e/ malonic
ation with
able sugge
):- Suggestedhloride/ureaid all forme
ee
-1) we ca
choline c
ination co
chemical
de. Additi
/urea effe
ment of two
n a weak l
s the conc
were weak
erties defe
d Chromiu
c acid al
h two malo
ested struc
d structuresa ionic liquied at room te
an see that
chloride/u
omplex w
series wh
ion of am
cted the
o urea mo
ligand fiel
centration
ker than N
ers in its c
um (III) c
lso at roo
onic acid
cture for C
s for Chromid with addemperature
63
t Chromiu
urea at ro
with six ur
here urea
mmonium
coordinat
olecules w
ld than ur
ns of urea
NO3-. This
coordinati
chloride h
om tempe
molecules
Cr (III) in
mium (III) ided salt NH4
e.
Re
um (III) c
oom temp
rea molec
considere
nitrate to
ion of Cr
with two N
rea accord
and chlo
s indicates
ion from
exahydrat
erature w
s as biden
both solv
in (1) cholin4NO3 and (3
esults an
chloride h
perature w
cule which
ed being i
the solut
r (III) wi
NO3- ions t
ding to the
oride are m
s that the
its free in
te in ionic
as indicat
ntate and tw
vents is sh
ne chloride/3) choline c
nd Discu
hexahydrat
was indic
h is in ag
in stronge
tion of Cr
ith urea m
took place
e spectroc
much hig
ionic liqu
ngredients
c liquid of
ted in oc
wo chlori
hown in Fi
/urea ionic hloride/ ma
ussion
te in the
cated in
greement
er ligand
r (III) in
molecule
e in spite
chemical
her than
uid has a
s in their
f choline
ctahedral
des. The
igure (3-
liquid, (2) alonic acid
Chapter Three Results and Discussion
64
Table (3-2):- The molar absorptivity of Chromium (III) ions in ionic liquid based on choline chloride.
Ionic Liquids Concentration
(M) [mol/l]
Concentration
(m) [mol/Kg]
Absorbance Molar absoptivity
(ε) [l cm-1mol-1]
Choline chloride/
Urea
0.1 M 7.4×10-2 m 0.406 4.06
Choline chloride/
Urea + NH4NO3
0.1 M 7.4×10-2 m 0.583 5.83
Choline chloride/
Malonic acid
0.1 M 7.5×10-2 m 0.037 0.37
3.2.2. Spectroscopic investigation of Iron (III) chloride solutions.
Iron (III) is isoelectronic with Manganese (II), but much less is known of the
greater tendency of the trivalent ion to have charge transfer bands in near
ultraviolet region with strong low energy wings in the visible that obscure the very
weak spin forbidden d-d bands. (64)
3.2.2.1. Iron (III) chloride in Choline chloride/Urea room temperature
ionic liquid solution.
Iron (III) chloride tetrahydrate was found to be completely soluble in choline
chloride/urea room temperature ionic liquid after 72 hours in concentration of
6.6×10-2m, producing initially a pale yellow solution which darkened to brownish
yellow color within the standing time of 72 hours. The pale yellow solution was
examined by ultraviolet visible and showed one absorption band at 15748 cm-1
(635.5 nm) and an absorption at the edge of ultraviolet absorption at 21881 cm-1
(457 nm), as illustrated in Figure (3-9), bands can be assigned to the following spin
forbidden transitions, the term symbol for the ground state of Fe (III) ion 6S can
split in tetrahedral crystal field as follows:- ν1=6 A1 → 4T1 and ν2= 6A1 → 4T2
Chap
Fig. (3-9)chloride/U
The
ultravio
Fig. (3-10Choline c
These
potassiu
The p
(III) com
complex
pter Thre
:- The U.V.-Urea ionic li
brownish
let absorp
0):- The U.Vchloride/Ure
e absorpti
um nitrate
pale yellow
mplex wit
xes conta
ee
-Vis. spectruiquid at roo
h yellow
ption at 21
V.-Vis. spectea ionic liqui
ions were
and Butyr
w solution
th chloride
aining ure
um of pale ym temperat
solution s
881 cm-1 (
trum of broid at room t
e close to
ramide -so
n might be
e ions (Fe
ea, or per
65
yellow solutture after 24
showed o
(457 nm),
ownish yellotemperature
o those fo
odium nitr
e raised d
eCl4-) whic
rhaps by
Re
tion of FeCl4 hours.
only the a
as shown
ow solution e after 72 ho
or Iron (I
rate,(67) as
due to the
ch might
more po
esults an
l3.4H2O of (6
absorption
n in Figure
of FeCl3.4Hours.
III) chlori
showed in
formation
be replace
lymerized
nd Discu
6.6×10-2m) i
n on the
e (3-10).
H2O of (6.6×
ide in ace
n Table (3
n of tetrah
ed by mor
d solvolyz
ussion
in Choline
edge of
×10-2m) in
etamide-
3-3).
hedral Fe
re stable
zed Iron
Chap
species.
electron
be sugg
two ure
chloride
melt aft
ultravio
3.2.2.2.
Addit
chloride
or in vi
spectros
the solu
was shi
visible d
Fig. (3-11addition o
This
therefor
pter Thre
Because
nic configu
ested that
ea and two
e in cholin
ter 40 day
let visible
Substituti
tion of
e/urea solu
iscosity o
scopy, and
ution, as sh
fted to hi
device cou
1):- The U.Vof NH4NO3 s
is not un
re no tran
ee
of that F
uration of
the envir
o chloride
ne chlorid
ys as ther
e spectrosc
ion React
ammonium
ution at ro
f the solu
d it indica
howed in
igher ener
uld not det
V.-Vis. specsalt room te
nexpected
nsitions ar
Fe3+ is an
f d5 with te
onment of
es in Td g
de /urea ro
re were n
copy.
tion with A
m nitrate
oom temp
ution. How
ated some
Figure (3
rgy of the
tect it.
ctrum of Feemperature
results as
e expecte
66
n oxidatio
erm symb
f the ligan
geometry.
oom tempe
no change
Ammoniu
e (NH4NO
erature di
wever, it
changes
3-11). The
e charge t
Cl3.4H2O ofionic liquid.
s the d5 c
ed under t
Re
on state w
ol of 6S th
nd for Fe3+
The prep
erature ion
es in its s
um Nitrate
O3) to t
id not sho
was exam
in the ultr
e edge ban
transfer re
f (6.6×10-2m.
configurat
the low li
esults an
with half-f
hat has no+ are of m
pared solu
nic liquid
spectra w
e.
the Iron
ow visual
mined by
raviolet vi
nd at 2188
egion whe
m) in Cholin
tion is no
igand field
nd Discu
filled d o
o splitting.
mixed mole
ution of I
d was stab
when exam
(III) in
changes i
ultraviole
isible spec
81 cm-1 (4
ere the ul
ne chloride/U
rmally sta
d as illust
ussion
orbital in
It could
ecules of
ron (III)
le in the
mined by
choline
n colour
et visible
ctrum of
457 nm),
ltraviolet
Urea with
able and
trated in
Chapter Three Results and Discussion
67
Tanabe-Sugano diagram Figure (3-13). Therefore it is not expected that the NO3-
and urea or chlorides are capable of inducing excitation to the Fe (III) ion to show
absorption in the visible region. However, a raised from the ultraviolet absorption
of the complex between Fe (III) and ligand as charge transfer spectra. Fe3+ is an
oxidation state with half-filled d orbital with term symbol of 6S that has no
splitting. It could be suggested that the environment of the ligand for Fe3+ are of
mixed molecules of two urea and two nitrates in tetrahedral geometry, when the
two chlorides were replaced with two nitrates molecules that is more stronger
ligand according to the spectrochemical series. The prepared solution of Iron (III)
chloride in choline chloride /urea with added nitrate room temperature ionic liquid
appeared to reach stable complex in the melt after 40 days, when it was examined
with ultraviolet visible spectroscopy no changes were observed.
3.2.2.3. Iron (III) chloride in Choline chloride/Malonic acid room temperature ionic liquid solution.
Iron (III) chloride tetrahydrate was found to be completely soluble in choline
chloride/ malonic acid room temperature ionic liquid after the standing time of 72
hours in concentration of 6.9×10-2m, producing a pale yellow solution. The pale
yellow solution was examined by ultraviolet visible spectroscopy and showed one
absorption band at 11255 cm-1 (888.5 nm) and an absorption close to an edge of
ultraviolet region at 18762 cm-1 (533 nm), and also showed another absorption at
14771 cm-1 (677 nm), as showed in Figure (3-12),the bands can be assigned as
shown below according to the term symbol for the ground state of Fe (III) ion 6S
for Td crystal field :- ν1=6 A1 → 4T1 , ν2= 6A1 → 4T2
Chap
Fig. (3-12ionic liqu
The
standing
coordina
Figure (
of a forb
lowest e
Fig. (3-1configur
pter Thre
2):- The U.Vid at room t
solution
g time of
ation of F
(3-13) the
bidden tra
energy.
13):- Tanaation.(64)
ee
V.-Vis. spectemperature
did not s
72 hours,
Fe (III) in
band that
ansition th
abe-Sugano
trum of FeCe.
showed an
, ultraviol
n this eu
t appeared
hat probab
o diagram
68
Cl3.4H2O of
ny chang
let visible
tectic. Us
d between
bly a raise
m for Oct
Re
f (6.9×10-2m
e in colo
spectra o
sing Tana
n the assig
d between
tahedral a
esults an
m) Choline c
or or in v
of the solu
abe-Sugan
gned band
n 6A1g and
and Tetrah
nd Discu
chloride/Ma
viscosity a
ution indic
no diagram
s was fou
d 4A1g, as
hedral for
ussion
alonic acid
after the
cates Td
m for d5
und to be
s it is the
r d5 ion
Chapter Three Results and Discussion
69
Because of Fe3+ is in oxidation state with half-filled d orbital with term symbol
of 6S that has so splitting. It could be suggested that the environment of the ligand
for Fe3+ are of mixed molecules of one malonic acid molecule and two chlorides in
tetrahedral geometry. The prepared solution of Iron (III) chloride in choline
chloride /malonic acid room temperature ionic liquid was stable in the melt after 40
days as there were no changes in its spectra when examined by ultraviolet visible
spectroscopy.
Iron (III) is the most stable oxidation state of the two oxidation states of Iron (Fe II-Fe III). Iron (III) in choline chloride/urea eutectic was found to be tetrahedraly
coordinated with chloride ions producing the pale yellow solution as FeCl4-, where
darken to brownish yellow because of the replacement of two chloride ions with
two more stable urea molecules. And that’s increased the absorbance of the visible
light in the blue region of energy as charge transfer band. Addition of ammonium
nitrate to the previous solution resulted in ligand substitution reaction between
chlorides and nitrate ions.
As illustrated in Figures (3-10), (3-11) and (3-12). It had been noticed that Fe3+
ions showed two absorption bands in choline chloride/ urea when examined after
24 hours (15748 cm-1 - 21881 cm-1) the first band disappeared leaving only the
higher energy band at 21881 cm-1 when examined after 72 hours. Further more,
when nitrate was added to the observe solution, the remaining band at 21881 cm-1
was also disappeared. This might be related to the slow reaction of Fe (III) with
solvent molecules in the first case attending equilibrium after 72 hours producing
more symmetrical coordination Iron (III) complex, the addition of nitrate ions
afford more stable d5 configuration producing no transition in the visible region,
and the absorptivities of the solutions were listed in Table (3-4).
In Table (3-3), Ultraviolet visible spectrum of Iron (III) in choline chloride/urea
eutectic showed two absorption bands initially, and only the edge band appeared
after 72 hours when the solution darkened. Addition of nitrate to the solution
Chapter Three Results and Discussion
70
shifted the edge absorbance band to higher energy charge transfer region where the
device showed it as a shoulder near the ultraviolet region. In choline chloride/
malonic acid eutectic the absorption bands were shifted to lower energies than that
in observed in water, and the absorbance in choline chloride/malonic acid was even
lower than in choline chloride/ urea eutectic Table (3-3). When malonic acid could
acts as a bidentate producing more stable Iron (III) complex while in choline
chloride/ urea eutectic the absorbance was higher because urea is monodentate
ligand and the formed Iron (III) complex was less stable. Figure (3-14) below
showed the suggested structures for Iron (III) in the prepared ionic liquids at room
temperature.
Table (3-3):- Electronic spectra of complexes of Iron (III) ions in various ionic liquids.
Ionic liquids Absorption band
(cm-1) Transitions Temperature
(oC) Suggested
structure Reference
Choline
chloride/urea
ν=15748 cm-1
ν = 21881 cm-1
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4T2 Room
temperature
Td
FeCl4- →[Fe Cl2
(urea)2]+
Present
work
Choline
chloride/urea
+ NH4NO3
About 457nm
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4T2 Room
temperature
Td
[Fe (NO3)2 (urea)2]+
Present
work
Choline
chloride/
malonic acid
ν1 =11255 cm-1
ν2=14771 cm-1
ν3=18762 cm-1
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4A1
(ν3) 6A1 → 4T2
Room
temperature
Td
[Fe (malonic acid)Cl2]-
present
work
Acetamide-
KNO3
ν=16900 cm-1
ν = 21000cm-1
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4T2 90oC Td
[FeCl4]- 67
Butyramide-
NaNO3
ν=16900 cm-1
ν =21000 cm-1
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4T2 90oC Td
FeCl4-
67
LiCl/KCl ν=5100 cm-1
6000 cm-1
(ν1) 6A1 → 4T1
(ν2) 6A1 → 4T2 400oC
1000oC
Td or
Distorted Td 1
LiNO3/KNO3 ν=1600 cm-1 160oC Td
[FeCl4]- 1
Chap
Table (3-4
Ionic Li
Choline
Urea
Choline
Urea + NH
Choline
Malonic ac
Fig. (3-14(2) cholinionic liqu
pter Thre
4):- The mo
iquids
C
chloride/ 0.
chloride/
H4NO3
0.
chloride/
cid
0.
4):- Suggestene chloride/ids all form
ee
lar absorpti
Concentration
[mol/l]
.1 M
.1 M
.1 M
ed structureurea ionic led at room t
ivity of Iron
n (M) Con
6.6×1
6.6×1
6.9×1
es of Iron (Iliquid with temperatur
71
n (III) ions in
centration (m
[mol/Kg]
10-2 m
10-2 m
10-2 m
III) complexadded NH4
e.
Re
n ionic liquid
m)
Abso
0.247
0.78
0.17
xes in (1) ch4NO3 salt, (
esults an
d based on c
orbance
holine chlori3) choline c
nd Discu
choline chlo
Molar abso
(ε)
[l cm-1m
2.47
7.8
1.7
ide/ urea iochloride/ma
ussion
oride.
optivity
mol-1]
nic liquid, alonic acid
Chap
3.2.3. S
3.2.3.1.
ionic liq
Coba
in cholin
72 hour
visual c
dissolut
The u
14903 c
(II) in ch
The a
consider
so in ord
from 6.9
in the sa
Fig. (3-15temperatu
pter Thre
Spectrosc
Cobalt
quid solu
lt (II) chlo
ne chlorid
rs in conce
changing
ion of the
ultraviolet
cm-1 (671
holine chl
absorbanc
red to be
der to mak
9×10-2m t
ame value
5):- The U.ure ionic liq
ee
copic inv
(II) chlo
ution.
oride hexa
de/ urea io
entration
in viscosi
salt.
t visible sp
nm) and 1
loride/ ure
ce of the
saturated
ke the spl
to 3.45×10
s, as show
V.-Vis specquid.
vestigatio
oride in
ahydrate w
nic liquid
of (6.9×1
ity of the
pectrum o
15600 cm-
ea eutectic
solution
when the
litting in th
0-2m, and
wn in Figu
ctrum of Co
72
on of Co
Choline
was found
d at room t
0-2m), pro
e choline
of the solu-1 (641 nm
c was cons
was high
splitting
he peaks m
the positi
ure (3-15).
oCl2.6H2O o
Re
balt (II)
chloride/
d to be eas
temperatur
oducing a
chloride/
ution show
m), as show
sistent with
her than
of the obs
more clea
ions of th
of (6.9×10-2
esults an
chloride
/Urea ro
sily and c
re after th
deep blue
urea ioni
wed two ab
wn in Figu
h tetrahed
expected
served pea
arly the so
he observe
2m) Choline
nd Discu
e solution
oom temp
completely
he standing
e solution
ic liquid a
bsorption
ure (3-18)
dral coordi
and the
aks was n
lution was
ed peaks r
e chloride/U
ussion
ns.
perature
y soluble
g time of
n without
after the
bands at
). Cobalt
ination.
solution
not clear,
s diluted
remained
Urea room
Chap
Coba
aqueous
tetrahed
coorden
Fig. (3-16Choline c
Using
and tetr
electron
electron
second t
band wa2G state
pter Thre
lt (II) chlo
s ionic li
dral coordi
nated as oc
6):- The U.chloride/Ure
g Orgel di
rahedral
nic config
nic transiti
transitions
as due to t
.
ee
oride show
iquid solv
inated com
ctahedral c
.V.-Vis. speea room tem
iagram fo
field Figu
guration th
ion of the
s were in t
the spin fo
wed a diff
vent than
mplex in th
complex w
ectrum of thmperature io
or the spli
ure (1-17
he appear
e tetrahed
the infrare
orbidden t
ν1=4
ν2=4
ν3=4
73
fferent beh
n in aque
he former
with 6 wat
he diluted nic liquid.
tting of fr
7) and Fi
red transi
dral comp
ed spectra
transition o
4A2 → 4T2
4A2 → 4T1
4A2 → 4T1
Re
haviour in
eous solv
r room tem
ter molecu
solution of
ree-ion 4F
igure (3-
ition was
plex of Co
a region. T
of the gro
2 (F),
1 (F)
1 (P)
esults an
n choline c
vent wher
mperature
ules, [Co(H
CoCl2.6H2O
F and 4P t
18) for t
assigned
o (II), wh
The splittin
ound state
nd Discu
chloride/ u
re it beh
ionic liqu
H2O)6]2+.
O of (3.45×
term in oc
the d7 tet
d to be t
here the f
ng in the a4A2 to the
ussion
urea non
haved as
uid and it
×10-2m) in
ctahedral
trahedral
the third
first and
appeared
e doublet
Chap
It was
chloride
the solu
tempera
amounts
showed
(512 nm
toward l
a change
As w
tetrahyd
of chlor
towards
Cobalt
absorpti
in the C
Cobalt
solution
liquid co
pter Thre
Fig. (3-17)
s obvious
e/ urea eut
ution, as it
ature ionic
s of calciu
octahedr
m) and wit
lower ene
e from oct
well as i
drate, the a
ride shift
s lower en
(II) nitrat
ion of the
Cobalt (II)
(II) form
n, the tem
omponent
ee
):- Crystal fie
that chan
tectic did
t reported
c liquids. I
um chlorid
al coordin
th addition
ergies and
tahedral to
n a mol
absorption
ted the ab
ergies and
te was oc
Cobalt (I
) species,
med compl
mperatures
ts as listed
eld splitting o
nging the c
not affect
d in other
In Calcium
de where t
nation wit
n of chlor
a large in
o tetrahedr
lten mixtu
n spectra o
bsorption
d pronoun
ccurred.(69
I) species
and also
lex(s) eff
or chang
d in Table
74
of 4F term of
concentrat
t the energ
studies o
m nitrate t
the solutio
th absorpt
ride cause
ncrease of
ral or dist
ure of 8
of Cobalt (
maximum
nced chang9) The eff
was attrib
in many
fected by
ging the m
(3-5).
Re
d7 ion Tetrah
tion of Co
gy of the u
of Cobalt
etrahydrat
on in pure
tion maxi
d a shift o
f the absor
orted octa
80% acetm
(II) chlorid
m of octa
ges in shap
fect of te
buted to s
other stud
changing
mole ratio
esults an
hedral config
obalt (II) c
ultraviolet
(II) chlori
te melts c
calcium n
imum bein
of the abs
rption coe
ahedral coo
mide-20%
de comple
ahedral C
pe of the i
emperature
structural c
dies where
g the con
o percent
nd Discu
uration.(64)
chloride in
t visible sp
ide in oth
ontaining
nitrate tetr
ng at 194
orption m
efficient in
ordination
% calcium
ex, where
Cobalt (II)
initial spec
e changes
changes o
e the geom
ncentration
of the us
ussion
n choline
pectra of
her room
variable
ahydrate
420 cm-1
maximum
ndicating
n.(68)
m nitrate
addition
) nitrate
ctrum of
s on the
occurring
metry of
n of the
sed ionic
Chap
Sugge
ionic liq
ligand in
The a
intense w
which i
tetrahed
shown i
6). The
eutectic
showed
solution
Figure (
solution
(533 nm
pter Thre
Fig. (3-18)
estion for
quid sugg
n tetrahed
absorbance
when com
indicates
dral or dis
in Table (3
solution
after 40
some cha
n (6.9×10-
(3-19), w
n showed t
m).
ee
:- Tanabe-S
the avera
gested that
dral coordi
e of the so
mpared wit
qualitativ
storted oct
3-5) and t
of Coba
0 days fro
anges in th-2m) as w
where the
two absorp
Sugano diagr
age enviro
t Cobalt (
ination, as
olution wa
th the sam
vely that
tahedral c
the molar
alt (II) ch
om the fi
he shape o
well as in
ultraviole
ption max
75
ram for d7 T
onment of
(II) were
s [CoCl42-]
as very lar
me concent
the coor
complex i
absorptiv
hloride hex
first ultrav
of the pea
diluted s
et visible
xima at 15
Re
Tetrahedral
f Cobalt (I
bonded w
].
rge and the
tration of
rdination
n choline
vity of the
xahydrate
violet visi
aks in the
solution (
spectra o
748 cm-1 (
esults an
ionic config
II) in chol
with four
e colour o
the salt in
of Cobal
chloride/
solution l
e in choli
ible spect
spectra of
(3.45×10-2
of concent
(635 nm)
nd Discu
guration.(64)
line chlori
chloride
of the solu
n aqueous
t (II) wa
/ urea eut
listed in T
ne chlorid
tra measu
f the conc2m) as sh
trated and
and at 187
ussion
ide/ urea
e ions as
ution was
solution
as either
tectic, as
Table (3-
de/ urea
urements
centrated
owed in
d diluted
761 cm-1
Chap
Fig. (3-19liquid aft(3.45×10-2
The u
spectra
molar a
formed
of the s
the Co (
The
substitu
which d
position
average
eutectic
molecul
chloride
than chl
(II) in c
substitu
pter Thre
9):- The U.Vter 40 days,2m).
ultraviolet
after and b
absoptivity
complex b
olution is
(II) in this
changing
tion of th
decrease t
ns are shi
environm
after 40
les in Cob
es with tw
loride acc
choline ch
ted by the
ee
V.-Vis. spect, (curve A)
t visible
before the
y of the s
become m
listed in
eutectic.
in the ult
e ligands
the energ
ifted to h
ment sugg
days face
balt (II) a
wo urea m
cording to
hloride/ ur
e stronger
tra of Cobalconcentrat
spectra o
e standing
solution d
more stable
Table (3-
traviolet v
that was
gy of the
higher ene
gested tha
ed some r
a complex
molecules,
o the elect
rea eutecti
urea mole
76
lt (II) in Ched solution
of the sol
time of 4
decreases
e and more
7) showed
visible spe
attached t
formed
ergies. Ca
at Cobalt
eplacemen
x. This s
which is
trochemica
ic formed
ecule.
Re
holine chlori(6.9×10-2m)
lution sho
40 days wh
after 40 d
e symmetr
d the effe
ectrum of
to the Cob
complex
alculation
(II) ions
nts of chl
ubstitution
preferable
al series.
d weak com
esults an
ide/Urea roo) and (curv
owed som
here the ab
days whic
rical. The
ct of the l
the solutio
balt (II) io
of Co (I
s using J
in cholin
lorides by
n took pl
e as urea
If so, it m
mplex wi
nd Discu
om temperave B) dilute
me differe
bsorbance
ch means
molar abs
ligand att
on indicat
on in the
II) and th
Jorgenson
ne chlorid
y more sta
lace betw
is stronge
means tha
th chlorid
ussion
ature ionic d solution
ences its
e and the
that the
soptivity
ached to
tes some
eutectic,
he peaks
rule of
de/ urea
able urea
ween two
er ligand
at Cobalt
de which
Chapter Three Results and Discussion
77
Table (3-5):- Electronic spectra for complexes of Cobalt (II) chloride in various room temperature
ionic liquids. Ionic liquids Percents of
IL.
Component.
Co (II) formed
complexes
Absorption
bands (cm-1)
Temp.
(o C). Suggested
structure.
References
Ca(NO3)2.4H2O-
CaCl2
Addition of
Cl-
[Co(NO3)4]2-
[Co(NO3)2Cl2]2-
[CoCl4]2-
ν=19420
Shifted to
lower energy
40-70 Oh
Distorted Oh
or Td
68
Acetamide-
Ca(NO3)2.4H2O
80%-20%
Addition of
Cl-
[Co(NO3)4]2-
[Co(NO3)2Cl2]2-
[CoCl4]2-
ν=19430
Shifted to
lower energy
40-90 Oh
Distorted Oh
or Td
69
KCl-AlCl3 (50.3-49.7)%
(50.3-49.7)%
(49.9-50.1)%
[CoCl4]2-
[CoCl6]4- or
Co(Al2Cl7)2
[Co(AlCl4)2]
_
25
300
300
Td
Oh
Distorted Oh
70
AICI3,-BPC 2:l
0.8:l
Co (Al2C I7) m2-m
(m=2)
CoCI42-
_
_
40
40
Oh
Td
71
butyramide-
Na NO3
(94.5-5.5)% [Co(butyramide)4Cl2] ν=1750
ν=8200
140 Oh 67
Acetamide-KNO3 [Co(acetamide)4Cl2] ν=1800
ν=8200
90-
140
Oh 67
LiCl-KCl [CoCl4]2-
ν=(14700,
15150,
16700)
447-
480
Td
1
3.2.3.2. Substitution Reaction with Ammonium Nitrate.
Addition of ammonium nitrate (NH4NO3) to the prepared solution at room
temperature did not show noticeable changes neither in color nor in viscosity of the
solution, but when it was examined by ultraviolet visible spectroscopy it showed
some changes in the ultraviolet visible spectrum of the Cobalt (II) ions, where the
shape and the position of the peaks changed as well as the absorbance, as shown in
Figure (3-20).
Chap
The u
(3.45×1
concentr
device c
bands w
(15552,
absorpti
to highe
formed
Fig. (3-20temperatu
The a
(P) and
state to
range ne
Calcu
replacem
Cobalt (
days, as
pter Thre
ultraviolet
0-2m) an
rated solu
could not d
with a clear
15848 cm
ion bands
er energy
complex o
0):- The U.Vure with add
appeared b
the splitti
doublet 2
ear infrare
ulations u
ment react
(II), no vi
s that hap
ee
visible sp
nd not th
ution sho
detect it. T
r splitting
m-1), (643,
positions
y indicated
of Cobalt
V.-Vis. spectdition of NH
band was a
ing in the
G state. W
ed region.
sing Jorg
tion took
isual obse
ppened to
pectrum o
he conce
owed uncl
The ultrav
in the ban
, 631 nm)
and the sp
d that ad
(II) in cho
trum of CoCH4NO3.
assigned t
e band wa
While the
enson Ru
place bet
erved chan
the origin
78
f the solut
entrated s
lear spect
violet visib
nd, the abs
) and 1890
plitting in
dded nitra
oline chlor
Cl2.6H2O of
to be of th
s due to f
second an
ule for av
tween add
nges happ
nal solutio
Re
tion was t
solution
tra with
ble spectru
sorption w
04 cm-1 (5
the band
ate made
ride/ urea
f (6.9×10-2m
he third tra
forbidden
nd first ba
verage env
ded nitrate
ened to th
on of Cob
esults an
taken to th
(6.9×10-2m
high abso
um showe
was at 150
529 nm), t
with smal
some int
melt.
m) in Choline
ansition fr
transition
and were
vironment
es and ch
he solution
balt (II) ch
nd Discu
he diluted
m), beca
orbance e
d three ab
015 cm-1 (6
the change
ll shift of t
eraction w
e chloride/U
rom 4A2 (F
n from 4A2
in the low
t suggeste
hlorides bo
n after tim
hloride in
ussion
solution
ause the
even the
bsorption
666 nm),
es in the
the band
with the
Urea room
F) → 4T1
2 ground
w energy
ed that a
onded to
me of 40
n choline
Chap
chloride
that took
From
two nitr
solution
stronger
Fig. (3-21temperatu
The
interelec
appropr
1117 cm
that for
more ex
tetrahed
pter Thre
e/urea ioni
k place to
m Figure (3
rates the
n as well a
r ligand th
1):- The U.Vure ionic liq
consisten
ctronic re
riate Tanab
m-1, the va
aqueous
xpansion
dral is less
ee
ic liquid a
the soluti
3-21), it is
complex
as molar a
han chlorid
V.-Vis. spectquid (curve A
ncy of th
epulsion R
be-Sugano
alue of B
solution
of d-elect
symmetri
at room te
ion after a
s obvious
became
absorptivi
de accordi
tra of CoClA) and with
his interp
Racah par
o diagram
was highe
(920 cm-1
tron charg
ical than o
79
emperature
ddition of
that after
less sym
ity increas
ing to the
l2.6H2O of ( addition of
pretation
rameter B
, Figure (3
er that for1) in octah
ge on com
octahedral
Re
e. Figure
f ammoniu
the replac
mmetrical,
ses; see T
spectroch
(3.45×10-2m)f NH4NO3 (c
have a
B, by fitt
3-18), and
r free ion
hedral coo
mplexation
l increases
esults an
(3-21) sh
um nitrate
cement of
as the a
Table (3-7)
emical ser
) in Cholinecurve B).
value o
ting the r
d it was fo
(971 cm-1
ordination
n in tetra
s the value
nd Discu
owed the
, as in tab
f two chlo
absorbance
), as the n
ries.
e chloride/U
obtained
ratio of ν
ound to be 1) and hig
n , becaus
ahedral as
e of B.
ussion
changes
le (3-6)
orides by
e of the
nitrate is
Urea room
for the
ν1/ν2 the
equal to
gher than
se of the
well as
Chapter Three Results and Discussion
80
3.2.3.3. Cobalt (II) chloride in Choline chloride/Malonic acid room
temperature ionic liquid solution.
Cobalt (II) chloride hexahydrate was found to be completely soluble in choline
chloride/ malonic acid ionic liquid at room temperature after the standing time of
72 hours in concentration of 7.4×10-2m, producing a blue colored solution without
any visual changing in viscosity of the choline chloride/ malonic acid ionic liquid
after the dissolution of the salt.
The ultraviolet visible spectrum of the solution showed three absorptions
maxima bands at 14881 cm-1 (672 nm), 15576 cm-1 (642 nm) and 16155 cm-1 (619
nm), that belong to the third transition band and the first and second transitions
were in the infrared and near infrared region respectively. The third band showed a
clear splitting and a spin forbidden transitions from 4A2 ground state to doublet
state with its splitting, indicating that Cobalt (II) chloride in choline chloride/
malonic acid ionic liquid at room temperature consist of Td coordination complex
of Cobalt (II), as shown in Figure (3-22), the term symbol for the ground state of
Co(II) ion 4F can split in tetrahedral crystal field, Using appropriate Orgal diagram
Figure (1-17), and Tanabe-Sugano diagram for the d7 tetrahedral electronic
configuration which is the same Tanabe-Sugano diagram for d3 octahedral
electronic configuration as in Figure (3-18), the transitions were assigned as
below:-
ν1= 4A2 (F) → 4T2 (F) =10Dq
ν2= 4A2 (F) → 4T1 (F)
ν3= 4A2 (F) → 4T1 (P)
The absorption bands of the solution were close to those of Cobalt (II) in choline
chloride/ urea room temperature ionic liquid in its initial readings after 72 hours
from the preparation, but Cobalt (II) showed changes in choline chloride/urea ionic
liquid where it did not show any changes in choline chloride/ malonic acid ionic
liquid, as showed in Table (3-6).
Chap
Fig. (3-22room tem
Calcu
Cobalt (
ligand i
any colo
shifts or
that Cob
that the
acid).
The
interelec
Tanabe-
the valu
aqueous
expansio
tetrahed
pter Thre
2):- The U.Vmperature io
ulations u
(II) in cho
n [CoCl42
or change
r changes,
balt (II) in
attached
consisten
ctronic rep
-Sugano d
ue of B w
s solution
on of d-
dral is less
ee
V.-Vis. spectrnic liquid.
using Jorg
oline chlor2-], and af
e or visco
, and the a
n choline
ligands w
ncy of th
pulsion pa
diagram, F
was highe
n 920 cm
-electron
symmetri
rum of CoC
genson Ru
ride/ malo
fter standin
osity chan
absorbance
chloride/
would not b
his interp
arameter B
Figure (3-1
er that for
m-1 in octa
charge o
ical than o
81
Cl2.6H2O of (
ule for a
nic acid e
ng time o
ge as wel
e of the so
malonic a
be replace
pretation
B , by fitti
18). And i
r free ion
ahedral c
on compl
octahedral
Re
(7.4×10-2m)
average en
eutectic wa
of 40 days
ll as abso
olution sta
acid eutec
ed by mor
have a
ing the ra
t was foun
n 971 cm-
oordinatio
lexation i
l increases
esults an
in Choline c
nvironmen
as bonded
s the solut
orption ban
ayed the s
ctic forme
re stable m
value o
atio of ν1/ν
nd to be eq-1 and hig
on , beca
in tetrahe
s the value
nd Discu
chloride/Ma
nt sugges
d with chlo
tion did n
nds did n
ame whic
ed stable c
molecule (
obtained
ν2 the app
qual to 11
gher than
ause of th
edral as
e of B.
ussion
alonic acid
sted that
orides as
not show
not show
ch means
complex,
(malonic
for the
propriate
117 cm-1,
that for
he more
well as
Chap
Fig. (3-23
choline ch
ionic liqu
Table (3-liquids.
Ionic li
Aqueou
solution
Chochlo/ ur
Cholinechloridurea +NH4NCholinechloridmalonicacid The energ
In the
that hav
with he
choline
pter Thre
3):- Suggeste
hloride/urea
id all forme
-6):- Electro
iquid Soluco
us
n
Pink
solut
oline orideea
Deepblue
e de/
NO3
Blue
e de/ c
Blue
solution fagies. e case of
ve been do
ating as s
chloride/
ee
ed structure
a ionic liqui
ed at room te
onic spectra
ution olor
Tranband
k
tion
ν1= 8
ν2= 1
ν3= 1
p ν1= ν2= ν3 =1560
e ν1= ν2= ν3 =1584
e ν1 = ν2 = ν3 =1615
aced some c
Cobalt (II
one on Co
showed in
/ urea an
es for Coba
id with add
emperature
a of comple
nsitions ds (cm-1) 8100
16000
19400
- -
= (14903-00) - -
= (15015- 48) - -
= (14881- 55) changes af
I) we saw
obalt (II)
n Table (3
nd cholin
82
alt (II) ions
ded salt NH
e.
exes of Cob
B (cm-1) 1
920 8
1117
1117
1117
fter 40 day
w that it be
in other i
-5) and T
ne chlorid
Re
in (1) choli
H4NO3 and (
balt (II) ion
0Dq sugg
100 Oh[Co
- Td [Co(NH
- Td[Co(
- Td[Co
ys and the
ehaved dif
ionic liqu
Table (3-6
de/ malon
esults an
ne chloride/
(3) choline c
ns in aqueo
gested struc
(H2O)6]+2
Cl4]2-→[CoCH2CONH2)2]
(NO3)2 (NH2C
Cl4]2-
readings h
fferently f
uids at roo
), where i
nic acid
nd Discu
/urea ionic
chloride/ma
us and var
cture Rn
Cl2 Pw
CONH2)2] Pw
Pw
ad shifted
from other
om temper
it was beh
ionic liq
ussion
liquid, (2)
alonic acid
rious ionic
Reference
64,65
Present work
Present work
Present work
to higher
r studies
rature or
haved in
quids as
Chapter Three Results and Discussion
83
tetrahedraly coordinated complex of [CoCl42-], that in choline chloride/urea ionic
liquid after a period of time up to 40 days, it faced some replacement of two
chlorides with two urea molecules, while in choline chloride/ malonic acid Cobalt
(II) did not faced any changes in its coordination with chlorides.
Table (3-7):- The molar absorptivity of Cobalt (II) ions in ionic liquid based on choline chloride.
Ionic Liquids
Concentration
(M) [mol/l]
Concentration
(m) [mol/Kg]
Absorbance
Molar
absorptivty (ε)
[l cm-1mol-1]
Choline chloride/
Urea
0.1 M 6.9×10-2 m 2.737 27.37
Choline chloride/
Urea + NH4NO3
0.1 M 6.9×10-2 m 2.022 20.22
Choline chloride/
Malonic acid
0.1 M 7.4×10-2 m 2.482 24.82
3.2.4. Spectroscopic investigation of Nickel (II) chloride solutions.
3.2.4.1. Nickel (II) chloride in Choline chloride/Urea room temperature
ionic liquid solution.
Nickel (II) chloride hexahydrate was found to be easily and completely soluble
in choline chloride/urea ionic liquid at room temperature after the standing time of
72 hours in 1.1×10-1m, producing a pale green solution without visual changing in
viscosity of the choline chloride/urea ionic liquid after the dissolution of the salt.
The ultraviolet visible spectrum of the solution showed three absorption maxima
the first one at 9522 cm-1(1050 nm), the second one at [13405 cm-1 (746 nm)-
14948 cm-1 (669 nm)] with clear splitting and the third one at 24390 cm-1 (410
nm), as shown in Figure (3-24), indicating that Nickel (II) in consistent with
octahedral coordination complex.
Chap
Fig. (3-24
temperatu
Using
term of
ground
While3A2g gro
pter Thre
4):- The U.V
ure ionic liq
g appropri
f d8 ion o
state of N
e the forbi
ound state
Fig. (3-25):
ee
V.-Vis. spect
quid.
iate Orgel
ctahedral
i (II) ion 3
idden tran
e to the sin
:- Crystal fie
trum of NiC
l diagram
configura3F can spli
ν1= 3
ν2= 3
ν3= 3
nsition was
nglet state
eld splitting
84
Cl2.6H2O of
Figure (1
ation Figu
it in octah
A2g→ 3T2
A2g→ 3T1
A2g→ 3T1
s due to sp
of 1Eg wh
of 3F term o
Re
f (1.1×10-1m
-17) and
ure (3-25)
hedral crys
2g (F)
1g (F)
1g (P) pin-orbit c
hich are ve
of d8 ion Oc
esults an
m) in choline
Crystal fi
. The term
stal field a
coupling t
ery close i
tahedral con
nd Discu
e chloride/ u
ield splitti
m symbol
as follows:
transition
in energy.
nfiguration.
ussion
urea room
ing of 3F
l for the
:-
from the
.
Chapter Three Results and Discussion
85
Using Tanabe-Sugano diagram for d8 electronic configuration Figure (3-26), the
third transition was recalculated and it was found to be similar to the value of the
experimental value, which is equal to 24390 cm-1 (410 nm). The value of ν1
represent the crystal field splitting value (10 Dq) indicating ligand field effect, the
different ligand field parameters have been calculated using Tanabe-Sugano
diagram for d8 electronic configuration.
The results of Nickel (II) in choline chloride/urea ionic liquid at room
temperature were close to those taken in aqueous solution, but it was not so similar
or close to those in other ionic liquid although Nickel (II) was found to have Oh
coordination, Table (3-8) show the results of Nickel (II) taken in aqueous solution
and some other ionic liquids at different temperatures.
Table (3-8):- Electronic spectra of complexes of Nickel (II) ions in aqueous solution and ionic
liquids.
Ionic liquid Ni (II) formed
complexes
Absorption bands
(cm-1)
Temp. (o C). Suggested
structure.
References
Aqueous
solution
[Ni(H2O)6]2+
ν1= 8500
ν2= 13800
ν3= 25300
Room
temperature
Oh
64,65,66
LiCl/KCl
[NiCl4]2-
ν1= 8000
ν2= 14200
ν3= 15300
700-1000
Distorted Td
1
CsCl
[NiCl4]2-
ν1= 7500
ν2= 14000
ν3= 15700
864
Td
1
Acetamide-
KNO3
[Ni(acetamide)4 Cl2]
ν1= 7500
ν2= 12500
ν3= 23200
90
Oh
67
Butyramide-
NaNO3
[Ni(butyramide)4 Cl2]
ν1= 8100
ν2= 12600
ν3= 23300
90-140
Oh
67
Chap
The c
the inter
ν1/ν2 on
to be eq
because
value of
This
delocali
only the
average
is small
in aqueo
Oh coor
Calcu
(II) in c
pter Thre
consistenc
relectronic
the appro
qual to 833
e of the le
f B.
phenome
ization of
e metal w
interelect
er (66), but
ous solutio
rdinated in
Fig. (3-26):-
ulation usi
choline ch
ee
cy of this i
c repulsio
opriate Tan
3 cm-1, the
ess expan
non is k
the meta
well. As a
tronic repu
t in cholin
on as well
n both cas
- Tanabe-Su
ing Jorgen
loride /ure
interpretat
n paramet
nabe-Suga
e value of
nsion of d
known as
l electron
result of
ulsion is r
ne chloride
l as for fre
es, as show
ugano diagr
nson Rule
ea eutecti
86
tion can b
ter ‘Racah
ano diagra
f B was lo
d-electron
s nephela
ns over mo
f this delo
reduced an
e/ urea eu
ee ion alth
wed in Ta
ram for d8 O
for averag
c is bonde
Re
e calculat
h’ paramet
am, Figure
ower than
charge on
auxetic ef
olecular o
ocalization
nd B (repr
tectic the
hough the
able (3-8).
Octahedral e
ge environ
ed with fo
esults an
ted and a v
ter B, by
e (3-26). B
that for fr
n comple
ffect and
orbitals th
n or “clou
resenting B
value of B
nickel wa
lectronic co
nment sug
our urea m
nd Discu
value obta
fitting the
B value w
ree ion 10
xation red
is attrib
hat encom
ud expand
B in the c
B was hig
as found to
nfiguration
ggested tha
molecules
ussion
ained for
e ratio of
as found
042 cm-1,
duce the
buted to
mpass not
ding” the
complex)
gher than
o be also
.
at Nickel
and two
Chap
chloride
stable ur
N
symmet
urea eu
smaller
molecul
chloride
Fig. (3-27
choline ch
3.2.4.2.
Addit
hexahyd
showed
The u
changes
the mol
pter Thre
es in octah
rea molec
Nickel (II
trical than
utectic is
see Table
les were b
es were bo
7):- The U.V
hloride/ urea
Substitu
tion of am
drate in c
visual cha
ultraviolet
s in the sh
lar absopt
ee
hedral coo
ule and ac
I) in chol
n in aqueo
smaller th
e (3-10), f
bonded to
onded with
V.-Vis. spec
a ionic liqui
ution Rea
mmonium
choline ch
anges neit
t visible sp
hape and t
tivity of
ordination,
ccording t
lnie chlor
ous solutio
han in aq
from the
Nickel (I
h Nickel (I
ctra of NiCl
id of (1.1×10
action with
nitrate (N
hloride/ ur
ther in the
pectrum o
the positio
the soluti
87
, where th
o the elect
ride/ urea
on, as the
queous so
absorbanc
II) ions w
II) with th
l2.6H2O in a
0-1m) at room
th Ammon
NH4NO3) t
rea ionic
e color nor
of the Nick
on of the
ion see T
Re
he chloride
trochemic
was fou
e absorban
olution an
ce we can
with in the
he z-axis, a
aqueous solu
m temperatu
nium Nit
to the solu
liquid at
r in the vis
kel (II) in
peaks and
Table (3-1
esults an
es were re
cal series.
und to be
nce in the
d the mo
n suggest
plane of
as in Figu
ution of 0.1
ure.
trate.
ution of N
room tem
scosity of
n this solut
d the abso
2), and t
nd Discu
eplaced w
more sta
choline c
olar absop
that the f
geometry
ure (3-32).
1 M (curve
Nickel (II)
mperature
the solutio
tion show
orbance as
the energi
ussion
ith more
able and
chloride/
ptivity is
four urea
y and the
A) and in
chloride
did not
on.
wed some
s well as
ies were
Chap
shifted t
and third
The u
the first
cm-1 (67
nm), ind
Nickel
bands.
Fig. (3-28NH4NO3 i
Using
field spl
While1Eg, due
forbidde
pter Thre
to lower e
d bands, a
ultraviolet
band at 9
77 nm), th
dicating th
(II) comp
8):- The U.ionic liquid
g appropri
litting wer
e the forb
e to spin-o
en transitio
ee
energy ban
as showed
t visible s
9522 cm-1
he third at
hat the ad
plex. The
.V.-Vis. speat room tem
iate Orgel
re assigned
3A
3A
3A
bidden tran
orbital co
on was fro
nds and th
d in Figure
pectrum o
(1050 nm
t 18149 c
dded nitrat
observed
ctrum of Nmperature.
diagram F
d as fallow
A2g (F) →
A2g (F) →
A2g (F) →
nsitions w
upling tha
om 3A2g to
88
he appeara
e (3-28).
of the solu
m), second
m-1 (551
te to the s
d energy b
NiCl2.6H2O
Figure (1-
ws, and sh
→ 3T2g (F)
→ 3T1g (F) (
→ 3T1g (P) (
were from
at mixes t
o singlet 1
Re
ance of new
ution show
band at 1
nm) and t
solution m
bands wer
in choline
-17) the tra
howed in F
(ν1) = 10D
(ν2)
(ν3) 3A2g grou
the 3T1g (F
A1g of the
esults an
w band be
wed four
3369 cm-1
the fourth
made som
re shifted
chloride/ u
ansition o
Figure (3-2
Dq
und state
F) and 1E
e 1G state.
nd Discu
etween the
absorption1 (748 nm)
h 24390 cm
me interact
d to lower
urea with ad
f Nickel in
25).
to singlet
Eg, and the
ussion
e second
n bands,
)- 14771
m-1 (410
ion with
r energy
ddition of
n crystal
t state of
e second
Chap
Tanab
re-check
value (1
Calcu
(II) in c
four ure
chloride
series th
urea is m
(II) in c
melt. Th
increase
pure me
Table (3
Fig. (3-29ionic liqu
A new
of Nick
reaction
pter Thre
be-Sugano
k the trans
10 Dq).
ulation usi
choline chl
ea molecu
es were re
his replace
more stab
choline ch
he solution
ed due to
elt was low
3-10).
9):- The U.Vid (curve A)
w transitio
kel (II) in
n that took
ee
o diagram
sitions and
ing Jorgen
loride /ure
ules and t
eplaced wi
ement is p
le and stro
loride/ ure
n of Nick
addition o
wer than i
V.-Vis. spectr) and pure m
on band a
choline c
k place be
Figure (3-
d the valu
nson Rule
ea eutectic
two nitrat
ith two nit
referable b
onger than
ea with ad
kel (II) in p
of NO3- an
in choline
ra of NiCl2.melt of choli
appeared in
chloride/
etween chl
89
-26) for d8
ue of (ν1) t
for averag
c with the
es in octa
trates, wh
because n
n nitrate. T
ddition of
pure melt
nd the mo
chloride/
6H2O in choine chloride/
n the ultra
urea eute
lorides an
Re
8 electroni
that refer
ge environ
e addition
ahedral co
hich accord
nitrates are
The absor
f nitrate m
t was less
olar absor
/ urea with
oline chlorid/ urea (curv
aviolet vis
ectic that
d added n
esults an
ic configu
to the cry
nment sug
of nitrate
oordinatio
ding to th
e stronger
rbance of
melt was hi
symmetri
rptivity of
h addition
de/urea withve B) at room
sible spect
pointed t
nitrates, as
nd Discu
uration was
ystal field
ggested tha
e was bond
on, where
e spectroc
than chlor
the soluti
igher than
ical as abs
f the Ni (I
n of NH4N
h addition om temperatu
tra of the
o the sub
s showed
ussion
s used to
splitting
at Nickel
ded with
the two
chemical
rides but
on of Ni
n in pure
sorbance
II) in the
NO3 as in
of NH4NO3 ure.
solution
bstitution
in Table
Chapter Three Results and Discussion
90
(3-9). This band was assigned to be in the value of 18149 cm-1(551 nm), and had a
probable a weak spin forbidden transition forbidden transition from triplet to
singlet as [3A2g (3P) → 1T1g (1G)], because it is the nearest forbidden transition
before the third triplet to triplet allowed transition due to mixing of spin-orbital
coupling.
The consistency of this interpretation can be calculated and a value obtained for
the interelectronic repulsion parameter or Racah parameter B, by fitting the ratio of
ν1/ν2 on appropriate Tanabe-Sugano diagram, Figure (3-26). And it was found to
be equal to 833 cm-1, the value of B was lower than that for free ion 1042 cm-1, and
close to the value B obtained in the pure melt 1082 cm-1, because of the less
expansion of d-electron charge on complexation reduce the value of B.
3.2.4.3. Nickel (II) chloride in Choline chloride/Malonic acid room
temperature ionic liquid solution.
Nickel (II) chloride hexahydrate was found to be easily and completely soluble
in choline chloride/ malonic acid ionic liquid at room temperature within the 72
hours in concentration of 1.2×10-1m, producing a bluish green solution without
visual changing in viscosity of the choline chloride/ malonic acid ionic liquid after
the dissolution of the salt.
The ultraviolet visible spectrum of the solution showed two absorption maxima
bands at 14184 cm-1 (705 nm) and 15220 cm-1 (657 nm) as show in Figure (3-30).
The ultraviolet visible spectrum indicated that Nickel (II) in choline chloride/
malonic acid eutectic consisted of Td complex or distorted Oh complex of Nickel
(II), but the Td complex is more preferable because it is closer to the value of
absorbance bands in aqueous solution of Nickel (II) chloride with high
concentration of chloride ions which indicates tetrahedral coordination as shown in
Table (3-9).
Chap
The tr
assigned
And the
transitio
transitio
transitio
Fig. (3-30
ionic liqu
Nick
when re
Orgel di
most lik
choline
pter Thre
ransitions
d as fallow
e appeare
on that oc
on from 3T
ons were in
0):- The U.V
id at room t
kel (II) for
eferred to
iagram d8
kely sugg
chloride/m
ee
of the N
w:
ed band i
ccurs in th
T1 ground
n infrared
V.-Vis. Spect
temperature
rmed Td c
Tanabe-S
Td and d
gestion is
malonic ac
Nickel (II)
3T1
3T1→3T1
in the sp
he visible
d state to s
d and near
trum of NiC
e.
complex i
Sugano di
d2 Oh have
that Nick
cid eutecti
91
in cholin
→3T2 (F)
→3A2 (F)
→3T1 (P)
ectra was
region w
singlet sta
infrared s
Cl2.6H2O of (
in choline
iagram of
e the same
kel (II) fo
ic.
Re
ne chlorid
(ν1)
(ν2)
(ν3)
s assigned
with a spli
ate of 1D.
pectra reg
(1.2×10-1m)
chloride/
f d2 Oh el
e transitio
ormed a T
esults an
de/ maloni
d to be t
itting due
While the
gion respec
in choline c
/ malonic
ectronic c
ons as in F
Tdcoordin
nd Discu
ic acid m
the third
e to the fo
e first and
ctively.
chloride/ ma
acid eutec
configurat
Figure (3-
nation com
ussion
melt were
allowed
orbidden
d second
alonic acid
ctic, and
tion, and
31), The
mplex in
Chap
Sugge
malonic
bonded
where t
accordin
malonic
transitio
containi
The a
choline
indicate
solution
complex
less sym
were sm
pter Thre
estion of
c acid wa
with one
the two
ng to the
c acid is
on with a s
ing high c
Fig. (3-31):-
absorbanc
chloride/
e that the
ns more st
x formed
mmetrical
maller in c
ee
the aver
as based o
e malonic
chlorides
electroch
stronger
split is sim
oncentrati
- Tanabe-Su
ce of the
/urea eute
Nickel (I
table and
in cholin
and less
choline ch
rage envir
on the sp
acid mol
were re
hemical s
than chlo
milar to th
ion of chlo
ugano diagr
solution
ectic befo
II) formed
more sym
ne chloride
stable com
hloride/ ur
92
ronment
ectra of t
lecule and
placed w
series this
orides. Th
hat resultan
oride ions
ram for d8 te
is higher
ore and a
d complex
mmetrical
e/malonic
mplex. Th
ea melt th
Re
of Nickel
the solutio
d two chl
with malon
s replacem
he observ
nt in aque
as in Tab
etrahedral e
than the
after the
xes in cho
l and defe
acid eute
he absorba
han in cho
esults an
l (II) in
on, where
lorides in
nic acid
ment is p
ved bands
ous soluti
ble (3-9).
lectronic co
solution
addition
oline chlo
erens in g
ectic, that
ance and
oline chlor
nd Discu
choline c
e Nickel
Td coord
molecule
referable
are of t
ion of NiC
nfiguration
of Nicke
of nitrate
oride/urea
geometry
t considere
molar abs
ride/ malo
ussion
chloride/
(II) was
dination,
e, which
because
the third
Cl2.6H2O
.
el (II) in
e, which
eutectic
than the
ed to be
soptivity
onic acid
Chapter Three Results and Discussion
93
melt. The tetrahedral complex of Nickel (II) d8 gives a term symbol of (3F), and the
transitions assigned as fallows:-
(ν1) = 3T1→3T2 (F), (ν2) = 3T1→3A2 (F), (ν3) = 3T1→3T1 (P)
The consistency of this interpretation can be calculated and a value was obtained
for the interelectronic repulsion parameter or Racah parameter B, by fitting the
ratio of ν1/ν2 from the appropriate Tanabe-Sugano diagram for d8 tetrahedral
configuration Figure (3-31), and it was found to be equal to 1030 cm-1, the value of
B was smaller than that for free ion (1042 cm-1) and higher for that in choline
chloride/ urea ionic liquid (833 cm-1), because in tetrahedral complexes the
expansion of d-electron charge is less than in octahedral complexes so the value of
B was reduced. Table (3-9) shows the results of Nickel (II) in choline chloride/
urea and choline chloride/ malonic acid ionic liquids at room temperature together
with those in aqueous solution.
As indicated in the Table (3-10), the absorption of Ni (II) ion had extended to
higher energy in choline chloride/ urea ionic liquid solvent in compared with those
observed in pure water. This is not unexpected, as urea is known to have stronger
field than water molecules and they are available large concentration in the melt.
However, when nitrate was added to Ni (II) solution in the previous melt solvent,
the absorption shifted even to higher energy than in pure melt. This is contrast to
the stronger field of urea than the nitrate as known in the spectrochemical series.
Therefore, the coordination of urea in pure melt with the metal ion indicated that
urea was already coordinated in the melt structure with other components, thereby
offering a reduced strength when coordinated with metal ions. Chloride ions
available in the melt compositions which are even weaker than urea and nitrate
ions in free state would play a negligible role in such coordination.
[I- < Br- <S2- < SCN- ~ Cl- <NO3- < N3
- < F- < CO(NH2)2 < OH- ~ ONO- <
CH2(COO)22- ≈ C2O4
2- < H2O < CN- ≈ CO].
Chapter Three Results and Discussion
94
Nickel (II) in the choline chloride/ urea and choline chloride/ malonic acid ionic
liquids solutions did not show any visual changes neither in color nor in viscosity,
and did not showed any spectral changes in the ultraviolet visible spectra of the
solutions after 40 days, which indicates stable complexes of Nickel (II). Figure (3-
32) showed the suggested structures for Nickel (II) in the prepared ionic liquids.
Table (3-9):- Electronic spectra of complexes of Nickel (II) ion in aqueous solution and various ionic liquids solution.
Solution Solution
color
Transitions
(cm-1)
B (cm-1) 10 Dq
(cm-1) Suggested
Structure
References
Aqueous solution Pale
green
ν1= 8500
ν2= 13800
ν3= 25300
1042 8500 Oh
[Ni(H2O)]2+
65
Aqueous solution
with high conc.
of chloride ions
Yellow ν1= -
ν2= 7549
ν3=(14250-
15240)
1030 4070 Td
[NiCl4]2-
66
Choline chloride/
Urea
Pale
green
ν1= 9522
ν2= (13405-
14948)
ν3= 24390
833 10000 Oh
[Ni(NH2CONH2)4Cl2]
Present
work
Choline chloride/
Urea +NH4NO3
Pale
green
ν1= 9522
ν2= (13369-
14771)
ν3= 24390
833 10000 Oh [Ni(NH2CONH2)4(NO3)2]
Present
work
Choline
chloride/Malonic
acid
Bluish
green
ν1= -
ν2= -
ν3cal.=(14184
-15220)
1030 - Td [Ni((CH2COOH)2)Cl2]2-
Present
work
Chap
Fig. (3-32choline chionic liqu
Table (3-1
Ionic L
Choline
Urea
Choline
Urea + NH
Choline
Malonic a
pter Thre
2):- Suggestehloride/ureaid all forme
10):- The m
Liquids
C
(
chloride/ 0
chloride/
H4NO3
0
chloride/
acid
0
ee
ed structurea ionic liquided at room te
olar absorp
Concentrati
(M) [mol/l]
0.1 M
0.1 M
0.1 M
es for Nicked with addeemperature
tivity of Nic
ion Co
(m)
1.1
1.1
1.2
95
el (II) ions ied salt of NHe.
ckel (II) ions
ncentration
) [mol/Kg]
×10-1 m
×10-1 m
×10-1 m
Re
in (1) cholinH4NO3 and (
s in ionic liq
n
Ab
0.054
0.178
0.844
esults an
ne chloride/ (3) choline c
quid based o
bsorbance
nd Discu
urea ionic chloride/ ma
n choline ch
Molar a
(ε)
[l cm-1m
0.54
1.78
8.44
ussion
liquid, (2) alonic acid
hloride.
absoptivity
mol-1]
Chap
3.2.5. S
Copp
from sin
characte
unsymm
easy to a
crystal f
of the s
33):-
3.2.5.1.
ionic liq
Copp
choline
hours in
visual c
dissolut
pter Thre
Spectrosc
per (II) com
ngle broad
erized by
metrical, b
assign una
field in the
spectra is
Copper
quid solu
per (II) chl
chloride/
n a concen
changing
ion of the
ee
copic inv
mpounds
d absorptio
large di
eing the r
ambiguou
e same wa
like wise
Fig. (3-33)
(II) chlo
ution.
loride dihy
urea ionic
ntration o
in viscosi
salt.
vestigatio
in genera
on band in
istortion f
result of a
usly. The f
ay as the te
expected
):- Crystal fi
oride in
ydrate wa
c liquid at
f 7.7×10-2
ity of the
96
on of Cop
al form blu
n the regio
from octa
number o
free ion gr
erm5D of
d(64) and ac
field splitting
Choline
as found to
t room tem2m, produ
e choline
Re
pper (II)
ue or gree
on (11000
ahedral sy
of transitio
round 2D t
the d4 ion
ccording t
g for 2D con
chloride/
o be comp
mperature
ucing a blu
chloride/
esults an
) chlorid
en colored
-16000) c
ymmetry
ons, which
term is exp
n and a sim
to the foll
nfiguration.
e/Urea ro
pletely and
after a st
ue colored
urea ioni
nd Discu
de solutio
d solution
cm-1. The
and the
h are by n
pected to
milar interp
lowing Fi
oom temp
d easily so
anding tim
d solution
ic liquid a
ussion
ons.
ns, raised
d9 ion is
bond is
no means
split in a
pretation
igure (3-
perature
oluble in
me of 72
n without
after the
Chap
The u
maxima
transitio
symbol
Fig. (3-34liquid at r
The c
same us
(II). Bec
using O
assign t
overlap
showed
transitio
Previ
Copper
the resu
pter Thre
ultraviolet
a for Cu (
on was as
with abro
4):- The U.Vroom tempe
color of th
sed salt fo
cause ther
Orgel diag
the value
of bands
a broad
ons.
ous studie
(II) behav
ults of Cop
ee
t visible s
(II) at 13
ssigned as
oad absorp
V.-Vis. Specerature.
he solution
or Cu (II)
re was no
gram Figu
of 10 Dq
. In the p
band at
es of Copp
ved differ
pper (II) in
spectrum
055 cm-1
s in Figur
ption band
ctrum of Cu
n was sim
which in
Tanabe-S
ure (1-16)
q, unfortu
present wo
(13055
per (II) in
rently from
n various s
97
of the sol
(766 nm
re (3-33) t
due to dis
uCl2.2H2O o
milar to the
ndicates an
Sugano dia
for d9 oc
unately, th
ork, the b
cm-1) wh
ionic liqu
m one sol
solvents.
Re
lution sho
m), as show
the only t
stortion of
of (7.7×10-2m
e color of
n octahed
agram for
ctahedral
his is mo
lue soluti
hich can
uids and aq
lution to o
esults an
owed one
wed in F
transition
f the Oh o
m) in cholin
the aqueo
dral coordi
the d9 ele
electronic
re difficu
on of Cop
be assign
queous sol
other, Tab
nd Discu
broad ab
igure (3-3
in d9 of
f Cu (II).
ne chloride/u
ous solutio
ination of
ectronic tr
c configur
ult becaus
pper (II)
ned to 2E
lution show
ble (3-11)
ussion
bsorption
34). The 2D term
urea ionic
on of the
f Copper
ransition,
ration to
e of the
complex
Eg→2T2g
wed that
showed
Chapter Three Results and Discussion
98
Copper (II) has d9 electronic configuration, it was suggested that the
environment of the ligands bonded with Cu2+ are of mixed ligands molecules of
four urea and two nitrates in Oh geometry.
Table (3-11):- The ultraviolet visible spectroscopic behavior of Copper (II) in various solvents.
Solvents Solution
color
Absorption
Band (cm-1)
Transition Suggested
structure
References
Aqueous solution Pale blue ν = (11000-
13000)
2Eg→2T2g Oh
[Cu(H2O)6]2+ 64,65
LiCl/KCl - ν = 9500 2B1g → 2B2g2B1g → 2Eg
Distorted Oh or
Td 1
AlCl3- 1,methyl
3ethylimidazoluim
chloride
- ν =34139
ν =24570
2B1g → 2B2g 2B1g → 2Eg
Td
[CuCl4]2- 72
Acetamide-KNO3 Blue-green
ν1 = 11000
(shoulder)
ν2 = 14000
ν3 = 20700
2Eg→2T2g
Oh
67
Choline chloride/
urea
Blue
ν1 =13055 2Eg→2T2g Oh
[Cu(CO(NH2)2)6]2+
Present
work
The Cu (II) solution in choline chloride/ urea eutectic did not show any visual
changes in color or viscosity as well as in the ultraviolet spectra after standing time
40 days, which means that Copper (II) already reached equilibrium and formed a
stable complex with urea molecules, in octahedral coordination.
3.2.5.2. Substitution Reaction with Ammonium Nitrate.
Addition of ammonium nitrate (NH4NO3) to the solution of Cu (II) in choline
chloride/ urea eutectic at room temperature did not show any visual changes
neither in color nor in viscosity of the solution, but it showed some changes in the
ultraviolet visible spectrum of the Copper (II) ions, as the absorbance and the
position of the peak had changed. The band was shifted to lower energy band at
average of 12903 cm-1(775nm) and a new band appeared on the edge of high
Chap
energy
interacti
Fig. (3-35addition o
The
absorban
interacte
transfer
in place
has a pr
is sever
so that t
Copp
that the
molecul
solution
was stab
examine
reached
two nit
pter Thre
level nea
ion betwee
5):- The U.Vof NH4NO3 a
absorbanc
nce than
ed Coppe
band. Th
ed in an e
rofound ef
rely distor
there is a p
per (II) is i
e environm
les of fou
n of Cu (II
ble in the
ed by ult
equilibriu
trates in
ee
ar charge
en added n
V.-Vis. specat room tem
ce of the
in pure m
er (II). Th
e d9 confi
nvironmen
ffect on al
rted, the ty
planar arra
ionic state
ment of
ur urea an
I) chloride
melt afte
traviolet v
um. The re
octahedra
e transfer
nitrate and
ctrum of Cmperature.
e solution
melt soluti
he new ab
iguration m
nt of cubi
ll its stereo
ypical dist
ay of four
e with d9 e
the ligan
nd two ni
e in cholin
er 40 days
visible sp
eplacemen
al coordin
99
at 2544
d Cu (II) i
CuCl2.2H2O
n when
ion which
bsorption
makes Cu
ic (i.e. reg
ochemistry
tortion is
short Cu-
electronic
ds bonde
itrates in
ne chloride
as there
pectroscop
nt took pla
nation tha
Re
45.3 cm-1
ion, as sho
in choline
nitrate w
indicates
band nea
u (II) subje
gular octa
y. When s
an elonga
-L bonds a
c configura
d with C
octahedra
e /urea roo
were no c
py, which
ace betwe
at tends t
esults an
(393 nm
own in Fig
chloride/ u
was added
s that the
ar the edg
ect to Jahn
ahedral) sy
six coordin
ation alon
and two tra
ation, it co
Cu2+ are o
al geomet
om tempe
changes in
h means t
en two ur
to be dis
nd Discu
m) indica
gure (3-35
rea ionic li
d showed
added nit
ge of high
n-Teller d
ymmetry,
nate “octa
g one 4-fo
ans long-o
ould be su
of mixed
try. The p
rature ion
n its spect
that the c
rea molecu
storted oc
ussion
ting the
).
iquid with
d higher
trate had
h charge
distortion
and this
ahedron”
fold axis,
ones.
uggested
ligands
prepared
nic liquid
tra when
complex
ules with
ctahedral
Chap
coordina
solution
symmet
results o
Fig. (3-36NH4NO3 i
3.2.5.3.
tempera
Copp
choline
of 72 h
without
liquid.
The u
broad b
Figure (
malonic
choline
(II) ions
pter Thre
ation. An
n from the
trical and
of the Cop
6):- The Uionic liquid
Copper
ature ion
per (II) chl
chloride/
hours in c
visual ch
ultraviolet
band was
(3-37), Th
c acid eut
chloride/
s was the s
ee
nd as the
e pure me
unstable,
pper (II) in
.V.-Vis. speat room tem
r (II) ch
nic liquid
loride dihy
malonic a
concentrat
hanging in
t visible sp
shifted to
he high ab
tectic mig
urea alon
same in al
absorban
elt solution
as show
n ionic liqu
ectra of Cumperature (c
hloride i
solution.
ydrate wa
acid ionic
tion of 8.
n viscosity
pectrum o
o lower e
bsorbance
ght indica
e or with
ll solutions
100
nce and m
n see Tab
wn in Figu
uids based
uCl2.6H2O icurve A), an
in Cholin
.
as found to
liquid at r
03×10-2m
y of pure
of the solu
nergy at
of solutio
ate a large
added nitr
s.
Re
molar abs
ble (3-13)
ure (3-36)
d on cholin
in choline cnd in pure m
ne chlor
o be easily
room temp
m, producin
choline c
ution show
9208 cm-
on of Copp
e distortio
rate althou
esults an
sorptivity
), the com
), Table (3
ne chlorid
chloride/ urmelt (curve B
ride/Malo
y and com
perature a
ng a deep
hloride/ m
wed one ab-1 (1086 n
per (II) in
on in its
ugh the co
nd Discu
increased
mplex beca
3-12) sho
de.
rea with adB).
onic acid
mpletely so
after stand
p yellow
malonic ac
bsorption
nm), as sh
n choline c
geometry
oncentratio
ussion
d in this
ame less
owed the
ddition of
d room
oluble in
ding time
solution
cid ionic
maxima
hown in
chloride/
than in
on of Cu
Chap
Fig. (3-37acid ionic
The r
those tak
one abs
Oh to T
choline
Copp
that the
molecul
Copper
one mal
molecul
acid is s
The s
show an
spectra
malonic
pter Thre
7):- The U.Vc liquid at ro
results of C
ken for Co
orption m
Td coordin
chloride b
per (II) is
e environm
les of one
(II) in ch
lonic acid
le is prefe
stronger th
solution of
ny visual
after it sta
c acid eute
ee
V.-Vis. spectroom temper
Copper (II
opper (II)
maxima at
nation. Ta
based prep
ionic stat
ment of
malonic a
holine chlo
d molecule
erable bec
han chlorid
f Copper (
changes
and for 40
ectic forme
rum of CuCature.
I) in choli
in molten
9500 cm-
able (3-12
pared ionic
te of d9 e
the ligan
acid and tw
oride/ mal
e. This rep
cause acc
de as ligan
(II) in cho
in the co
0 days, wh
ed stable c
101
Cl2.6H2O of (
ine chlorid
n chloride -1, where C
2) showed
c liquids.
lectronic
ds bonde
wo chlorid
lonic acid
placement
ording to
nd.
oline chlor
olor or vi
hich means
complex.
Re
(8.03×10-2m)
de/ maloni
of LiCl/K
Copper (I
d all the r
configura
d with C
des in dist
was bond
t of two c
the spec
ride/ malo
iscosity as
s that Cop
esults an
) in choline c
ic acid eu
KCl at 400
II) prosed
results tak
ation, it co
Cu2+ are o
orted Oh o
ded with t
chlorides w
trochemic
onic acidio
s well as
pper (II) in
nd Discu
chloride/ ma
utectic is si
0oC which
to be in d
ken for Cu
ould be su
of mixed
or Td geom
two chlor
with malo
cal series
onic liquid
in the el
n choline c
ussion
alonic
imilar to
h showed
distorted
u (II) in
uggested
ligands
metry.
rides and
onic acid
malonic
d did not
lectronic
chloride/
Chapter Three Results and Discussion
102
Table (3-12):- Electronic spectra of complexes of Copper (II) ions in choline based ionic liquids at
room temperature.
Solution Solution
color
Absorption
band (cm-1)
Transitions
Suggested structure References
Choline
chloride/ urea
Blue
ν = 13055 2Eg→2T2g Oh
[Cu(CO(NH2)6]2+
Present work
Choline
chloride/ urea
+NH4NO3
Blue
ν =12903 2Eg→2T2g Oh
[Cu(CO(NH2)4(NO3)2]
Present work
Choline
chloride/
malonic acid
Yellow
ν =9208 2B1 → 2B2 2B1 → 2E
Distorted Oh or Td
[Cu((CH2COOH)2)Cl2]2-
Present work
Table (3-13):- The molar absorptivity of Copper (II) ions in ionic liquid based on choline chloride.
Ionic Liquids
Concentration
(M) [mol/l]
Concentration
(m) [mol/Kg]
Absorbance
Molar absoptivity
(ε)
[l cm-1mol-1]
Choline chloride/
Urea
0.1 M 7.7×10-2 m 1.702 17.02
Choline chloride/
Urea + NH4NO3
0.1 M 7.7×10-2 m 2.005 20.02
Choline chloride/
Malonic acid
0.1 M 8.03×10-2 m 2.312 23.12
Chap
Fig. (3-38choline chionic liqu
pter Thre
8):- Suggestehloride/ureaid all forme
ee
ed structurea ionic liquided at room te
es for Copped with addeemperature
103
er (II) ions ed salt of NHe.
Re
in (1) cholinH4NO3 and
esults an
ne chloride/(3) choline
nd Discu
/ urea ionic chloride/ma
ussion
liquid, (2) alonic acid
Chapter Three Results and Discussion
104
3.3. Conclusion.
1) The prepared room temperature ionic liquids was based on quaternary
ammonium salt (choline chloride) that was mixed with hydrogen bond donor
molecules like urea [CO(NH2)2] in mole ratio of 1:2, and malonic acid
[(CH2COOH)2] in mole ratio of 1:1. Showed different physical properties
from other non aqueous solvent.
2) Those ionic liquids have been used as non-aqueous solvents to study the
behavior of some first row transition metal compounds. The interaction that
took place between transition metal ions and the possible ligands in each
solution gave complexes comparable geometry.
3) All the prepared solutions of the transition metal ions were colored solution,
while the prepared ionic liquids were colorless.
4) The new formed complexes solution of each transition metal ion in each ionic
liquid was analyzed by ultraviolet visible spectroscopy, and the spectra of the
solutions were studied.
5) The prepared ionic liquids showed different ligands field strength properties,
because ionic liquids contents had molten properties.
6) The ligand field calculated showed different strength (according to the 10 Dq
value), different Racah parameter (B) and different electron- repulsion
parameter (β), which refers to different ionic character between the metal and
donor atoms of the ligands.
7) The infrared spectroscopy of the solutions of transition metals in ionic liquids
was difficult to establish because the ionic liquids were insoluble in the
solvents were used in the infrared spectroscopy instrument and it was soluble
only in absolute ethanol, as well as it was not able to use the ionic liquids
without dilution because the solution will be over saturated for the IR chart
and the peaks will not be clear.
Chapter Three Results and Discussion
105
8) The conductivity of the transition metals in ionic liquids solutions was also
difficult to determine because the ionic liquids solution is conductance
himself, and it was hard to extract the formed complex of transition metal
with ionic liquids from its solution with the known extraction method, to
determine the conductivity of the formed complex.
9) The magnetic susceptibility of transition metal in ionic liquids solutions was
difficult to calculate because the solutions were thick and hard to force into
the magnetic susceptibility tubes with injection because of the formation of
air bobble in the bottom on the tube that couldn’t be gotten red of it by
shaking or worming the tube.
Chapter Three Results and Discussion
106
3.4. Suggestion for future work.
The following suggestions can be postulated to ensure completion of the study
in this work.
1) Preparation of another set of transition metal complexes using other transition
metal ions such as second or third rows of transition metal using the same
prepared ionic liquids.
2) Preparation of another set of transition metal complexes using other reported
room temperature ionic liquids or newly prepared ionic liquids.
3) Preparation of the same solutions with the same concentration of transition
metal ions and by using the same analyzing instrument but in various
temperature ranges.
4) Preparation of the same solutions of transition metal ions using different
concentration of the used transition metal compounds.
5) To complete the measurements and identification of the new formed
complexes of transition metal ions by using infrared spectroscopy, magnetic
susceptibility measurements and x-ray diffraction.
6) Find a suitable method to isolate the prepared complexes of transition metal
ions from their ionic liquid solution.
References
107
1. J. J. Lagowski, “The chemistry of non-aqueous solvents”, vol. 5B, New York,
1978.
2. J. E. Huheey, “Inorganic Chemistry, Principles and reactivity”, Harper, New
York, 1972.
3. Bruce N. Hoglund, “What is Molten Salts and Its Technology”, 1997.
4. Alvin Martin Weinberg, “The first Nuclear Era; the life and times of a
technological fixer”, page 198-200, 1994.
5. H. G. Mac Pherson, “The Molten Salts Adventure”, vol. 90, page 374-380,
1985.
6. Tames A. Lane, H.G. Mac Pherson and Frank Maslan, “Fluid Fuel Reactor”,
1958.
7. R. C. Briant et al., “Nuclear Science and Engineering”, vol. 2, no. 6, page
795-853, 1957.
8. Temma Vickery, “Green chemistry”, University of Bristol, 2001.
9. Anastas P. T., Warner J. C., “Green Chemistry Theory and Practice”, 1999.
10. Cann M.C. and Connelly M.E., “Real world cases in green chemistry”;
American Chemical Society, 2000.
11. G.P. Smith, A.S. Dworkin, R.M. Pabni, and S.P. Zingg; J. Am. Chem. Soc.,
vol. 3, page 525, 1989.
12. Di chio Donata, Fusco Marida, Halo Migen, “Ionic liquids”, University of
Torino, 2004.
13. Martyn J. Earle and Kenneth R. Seddon, “Ionic liquid, Green solvents for the
future.”; Pure Appl. Chem., vol. 72 , no. 7, page 1391-1398 , 2000.
14. M. Freemantle; Chem. Eng. News, vol. 76, (30th march) page 32-37, 1998.
15. T. Welton; Chem. Rev., vol. 99, page 2071-2083, 1999.
16. J. Holbrey and K.R. Seddon; Clean Prod. Proc., vol. 1, page 223- 236, 1999.
17. K.R. Seddon; J. Chem. Tech. Biotech., vol. 68, page 351-356, 1997.
18. H. L. Chem, V. R. Kock, L. L. Miller, R.A. Osteryoung; Jor. Am. Chem.
Soc., vol. 97, page 3264- 3267, 1975.
References
108
19. P. Walden; Bull. Acad. Imper; Sci., (St. Petersburg), page 1800, 1992.
20. J.S. Wilkes, J.A. Levisky, R.A. Wilson, C.L. Hussey; Inorg. Chem., vol. 21,
page 1236- 1264, 1982.
21. K.R. Seddon., “In molten salt forum; proceedings of 5th international
conference on molten salt chemistry and technology”, vol. 5-6, H. Wendt ed.,
page 53- 62, 1998.
22. C.L. Hussey, T.B. Scheffler, J.S. Wilker, A.A. Fannin; J. Electro-Chem. Soc.,
vol. 133, page 1389-1391, 1986.
23. J.D. Holbrey and K.R. Seddon; J. Chem. Soc., Dalton Truns, page 2133-
2139, 1999.
24. C.M. Gordon, T.D. Hollrey, A.R. Kennedy, K.R. Seddon, J. Mater ; Chem.,
vol. 8, page 2627-2636, 1998.
25. C. Hardacre, K.R. Seddon, J. McGrath, “Room temperature Ionic liquids”,
Queens University Belfast, 2000.
26. Hiroyuk; Ohno, “Electro Chemical Aspects of Ionic liquids, chapter one: -
Importance and possibility of Ionic liquids”, John Wiley and Sons, Inc., 2005.
27. John S. Wilakes, “A short history of liquids– from molten salts to neoteric
solvents”; Green Chemistry, The Royal Society of Chemistry, vol. 4, page 73-
80, 2002.
28. J.L. Copland, “Transport properties in Ionic liquid”, Gordon and Breach
science publishers, New York, 1974.
29. C. L. Hussey, “The Electro Chemistry of room-temperature haloaluminate
molten salts”, (ed.) G. Mamantov and A.I. Popov, “Chemistry of non-aqueous
solutions”, VCH publisher Inc. New York, page 227- 276, 1994.
30. R.T. Carlinand, J.S. Wlikes, “Chemistry and speciotion in room-temperature
chloroaluminate molten salts”, (ed.) G. Mamantov and A.I. Popov,
“Chemistry of non-aqueous solutions”, VCH publisher Inc., New York, page
277-307, 1994.
References
109
31. R.T. Adams, Ger. Pat., 3347947, 1967; G.P. Hugen and D.T. Hung, Br. Pat.
2246788, 1992; E. Ceausescu, M. Corciovei, T. Sarbu, A. Schramm and D.
Stadermann, Plaste Kautsch., vol. 35, page 378, 1988.
32. Jairton Dupont, Crestina S. Consorti and John Spencer, “Room Temperature
Molten Salts: Neoteric “Green” Solvents for Chemical Reactions and
Processes”; J. Braz. Chem. Soc., vol. 2, no. 4, page 337- 344, 2000.
33. Bolm C., Beckmann O., Dabard O. A. G.; Angew. Chem., vol. 3, page 957,
1999; Angew. Chem. Int. ed., vol. 38, page 907, 1999.
34. Metzger J.O.; Angew. Chem., vol.110, page 3141, 1998; Angew. Chem. Int.
ed., vol. 37, page 2975, 1998.
35. Jessop, P.G. I. Kariya, T. Noyori R.; Chem. Rev., vol. 99, page 475, 1999.
36. Walden, P. Bull, Acud. Imper. Sco. (St. Petersburg), page 405. C.A.1914, vol.
8, page 2291, 1914.
37. C.L. Hussey, in “Advances in molten salts chemistry”; Mamantov G.;
Mamantov C. (eds.); Elsevier , New York , vol. 3, page 185, 1983.
38. Suarez, P. A. Z., Einloft S., Dullius, J.E.L., de Souza, R.F.; Dipont. J. Tor.
Chem. phys., vol. 95, Page 1626, 1998.
39. Frank Endres and Sherif Zien El Abedin, “Air and water stable ionic liquid in
physical chemistry”; Phys. Chem. Chem. Phys., vol. 8, no. 18, Page 2101-
2116, 2006.
40. H.L. Chum, V.R. Koch, L.L. Miller and R.A. Osteryoung.; J. Am. Chem.
Soc., vol. 97, Page 3264, 1975.
41. J. Robinson and R.A. Osteryoung.; J. Am. Chem. Soc., vol. 101, Page 323,
1979.
42. J.S. Wilkes, J.A. Levisky, R.A. Wilson and C.L. Hussey; Inorg. Chem., vol.
22, page 2099, 1983.
43. T.B. Scheeffler, C.L. Hussey, K.R. Seddon, C.M. Kearad, P.D. Armitage;
Inorg. Chem., vol. 22, Page 2099, 1983.
References
110
44. D. Appleby, C.L. Hussey, K.R. Sedden and J.E. Trup; Nature, vol. 323, page
614, 1986.
45. C. L. Hussey; “Advanced Molten Salt Chemistry”, vol. 5, page 185, 1983.
46. P.C. Trulove and R.A. Mantez, “In Ionic liquid in Synthesis”, (ed.) P.
Wasserscheid and T. Welton, Wiley –VCH, Weinheim, Page 103-126, 2003.
47. R.A. Mantez and P.C. Trulove, “In Ionic liquids, in synthesis”, (ed.) P.
Wasserscheid and T. Welton, Wiley –VCH, Weinheim, Page 56-68, 2003.
48. J.G. Huddleston, A.E. Visser, W.M. Reichert, H.D. Willauer, G.A. Broker
and R.D. Rogers; Green Chem., vol. 3, Page 156, 2001.
49. P. Wasserscheid and W. Kiem; Angew Chem. Int. ed., vol. 39, Page 3772,
2000.
50. J. Fuller, R.T. Carlin and R.A. Osteryoung; J. Electro Chem. Soc., vol. 144,
Page 3881, 1997.
51. J.D. Hlbrey and R.D. Rogers, “In Ionic liquid in synthesis”, (ed.) P.
Wasserscheid and T. Welton, Wiley –VCH, Weinheim, Page 41-55, 2003.
52. M.E. Van Valkenburg, R.L. Vaughn; Thermo Physical Properties, 2003.
53. J. Sun, M. Forsyth and D.R. Mac Farlane, “Room Temperature Molten Salts
Based on the Quaternary Ammonium Ion.”; J. Phys. Chem., B, vol. 102, Page
8858-8864, 1998.
54. UNEP Publications, Foreword Introduction of Choline Chloride, Berlin,
Germany, 2000.
55. Heck, R.F., “Organic Reaction”; John Wiley and Sons, New York, vol. 27,
Page 345, 1982.
56. Chang Biying, Song Rong, Ding Yong Sheng and Liew Chee Yong,
“Determination of Choline chloride content”, Asian Poultry Magazine, Aug.
2002.
57. P. Wasserscheid and T. Welton, “Ionic liquid in synthesis”, Weinheim, Wily-
VCH Verlag, 2003.
References
111
58. Andrew P. Abbott, Glen Capper, David L. Davies, Raymond K. Rasheed and
Vasuki Tambyrajah, “Novel solvent properties of choline chloride/urea
mixture”; Chem. Commun., page 70-71, 2003.
59. M.S. Sitze, E.R. Schreiter, E.V. Patterson and R.G. Freeman; Inorg. Chem.,
vol.40, page 2298, 2001.
60. L.Z. Liu, Y.X. Tong and Q.Q. Yang; Rare Met, vol. 19, page 237, 2000.
61. A.P. Abbott, G. Capper, D.L. Davies, H. Munro, R.K. Rasheed and V.
Tambyrajah; Chem. Commun., page 2010-2011, 2002.
62. Andrew P. Abbott, David Boothdy, Glen Capper, David L. Davis and
Raymond K. Rasheed, “Deep eutectic solvent formed between choline
chloride and carboxylic acids”; J. Am. Chem. Soc., vol. 126, no.29, page
9142-9147, 2004.
63. Andrew P. Abbott, David Boothdy, Glen Capper, David L. Davis and
Raymond K. Rasheed, “Ionic Liquids and Their Use as Solvents”, 2000.
64. Frank Albert Cotton and Geoffrey Wilkinson; Advanced Inorganic chemistry,
4th edition, A Wiley intercince publication, 1980.
65. James E. Huheey; Inorganic chemistry: principles of structure and reactivity,
Harper international, S.I. edition, 1983.
66. N.N. Greenwood and A. Earnshaw; Chemistry of the elements, A Wiley
intercince publication, 3rd edition, 1986.
67. E.I. Eweka and D.H. Kerridge, “Solution chemistry of molten amide- nitrate
eutectic”; Chem. Papers, vol.53, no.1, page 11-15, 1999.
68. Istvan J. Zsigrai, Slobodan B. Gadzuric and Ruzica Nikolic, “Cobalt (II)
chloride complexes in Calcium nitrate tetrahydrate melt”, Institute of nuclear
sciences ‘Vinca’, 2000.
69. Jelena Savovie, Ruzica Nikolie, Agan Veselinovie, “Cobalt (II) chloride
complex formation in Acetamide- Calcium nitrate tetrahydrate melt”; J. of
solution chemistry, vol. 33, no. 3, page 287-300, March 2004.
References
112
70. Harald A. Øye and Deiter M. Gruen, “Cobalt (II) species in fused chloride
solvents”; Inorg. Chem., vol.4, no. 8, Aug. 1965.
71. C.L. Hussey and T.M. Laher, “Electrochemical and spectroscopic studies of
Cobalt (II) in molten aluminium chloride N,n-butylpyridium chloride”; Inorg.
Chem., vol.20, no. 12, page 4201-4206, 1981.
72. T.M. Laher and C.L. Hussey, “Copper (I) and Copper (II) cholro complexes
in the basic aluminium chloride 1-methyl, 3-ethylimidazolium chloride, ionic
liquid”; Inorg. Chem., vol.22, no. 22, page 3247-3251, 1983.
73. R. J. Lancashire, Chemistry of the First Row Transition Metal Complexes,
Inorganic Chemistry, September 2005.
74. Calculations using Orgel Diagram and Tanabe-Sugano diagram, From
Wikipedia, the free encyclopedia (www.wikipedia.org), 2005.
75. CHEM 340, Lect-14, the using of Tanabe-Sugano diagram. (www.
encyclopedia.org), 2004.
جمهورية العراق وزارة التعليم العالي والبحث العلمي
جامعة النهرين كلية العلوم
قسم الكيمياء
العناصر بعض مالحأل طيفيتشخيص األمالح منصهرات يفاألنتقالية
درجة حرارة الغرفةب
رسالة جامعة النهرين-مقدمة الى كلية العلوم
ء في الكيمياوهي جزء من متطلبات نيل درجة الماجستير
من قبل
مريم علي سعيد ٢٠٠٥ )جامعة النهرين(بكالوريوس
بأشراف
الدكتور هادي محمد علي عبود
۲۰۰٨ أيار ۱٤۲٩ ربيع الثاني
حيمبن الرمحم الله الرس أحدني رأيتإأبت ـ ي ألبيهسفذ قـال يوإ
ي ـ ل والقـمرلشمس كـوكبا و اعشرءياك بني ال تقصص رـقـال ي۞جدينـس
ن إ آك كيدـخوتك فيكيدوا لإعلى و كذلك ۞بينـ م˝سن عدوـن لألنـالشيط
من تأويل األحاديث بك ويعلمكيجتبيك ر عليك وعلى آل يعقوب كما أتمها ويتم نعمته
ن إق ـسحإ و براهيمإعلى أبويك من قبل ۞ حكيمربك عليم
ظيماهللا الع قدص
الذي أحببتفيا والدي أهديكم ذا األثرن كنتٳ
جميال الجميلردأفأني بنور الضمير ونور البصر
بروحي فأن كان بحثي سماءاو القمرفأنتم شمسها
وكل من له ...أهلي و أصدقائي...اخي و أختي...والدي العزيز...الى والدتي الغالية
....مكان في قلبي
...ور أياد حمزة جاسم الخفاجيالى روح أستاذي العزيز األستاذ الدكت
.اهدي عملي المتواضع
مريم علي سعيد
األهداء
آمذيب نوعين من األمالح المنصهرة في درجة حرارة الغرفة لألستخدامها في هذا البحث تم تحضير
تجاه أيونات المرآبات األيونية السائلة المائي في دراسة الهدف منها التعرف على السلوك الذي تسلكه هذه
العناصر األنتقالية وهذه المرآبات حضرت من مزج ملح مشتق من جذر األمونيم الرباعي و مواد لها قابلية
بنسب مولية ) اليوريا و حامض آاربوآسيلي ثنائي الكاربوآسيل(على تكوين أواصر هيدروجينية مثل
.ثابته
دات الع ة بعض معق ة تمت دراس ي درج صهرة ف الح المن ن األم د م وع جدي ي ن ة ف ناصر األنتقالي
.حيث أن األمالح المنصهرة تعتمد في ترآيبها على آلوريد الكولين, حرارة الغرفة
دة دات الجدي آ شخصت المحاليل المحضرة التي تحتوي على المعق صآ طيفي وق تشخي بواسطة األشعة الف
سجية ة -البنف ث ل). UV-Vis(المرئي الح حي ات األم الي و مكون ون األنتق ين األي دات ب ون معق وحظ تك
ة د و . المنصهرة في درجة حرارة الغرف ولين آلوراي ا و آ د يوري ولين آلوراي األمالح المنصهرة سميت آ
. درجة مئوية على التوالي٠ و ١٠حامض المالونك و الذان يمتلكان درجة أنجماد في
ن ل م سطوح لك ة ال كال ثماني ت ذات أش ا آان د يوري ولين آلوراي ول الك ي محل ات ف وحظ أن األيون ل
ات +Cu٢ ,+Ni٢ ,+Cr٣أيونات حيث أظهرت دراسة . +Co٢ ,+Fe٣ و أشكال رباعية السطوح لكل من أيون
ي األنتقاالتالطيف دات ثماني dn المسموحة و الممنوعة لكل ترآيب ألكترون سطوح و في آل من المعق ة ال
.رباعية السطوح
ا د يوري ولين آلوراي ول الك الي في محل تمت دراسة تفاعل أستبدال لليكندات المرتبطة باأليون األنتق
ول و , بأضافة أيون النترات على شكل نترات األمونيوم ات المحل ين جزيئ اط ب حيث أظهرت ضعف األرتب
قية كال تناس ي أش الي الموجب ف ون األنتق ل المع. األي كال آ ي أش ستقرة ف االت م رت ح ة أظه دات المتكون ق
د , ساعة من تحضيرها ٧٢معقداتها خالل ه بع ر في طيف عدا في حالة أيون الكوبلت الثنائي حيث أظهر تغي
. يوم مما يدل على عدم وصول المعقد المتكون حالة األستقرار٤٠
سطوح +Cu٢ ,+Ni٢ أما في محلول الكولين آلورايد و حامض المالونك فأن أيوني أظهرا شكل رباعي ال
.أآثر من أشكال ثمانية السطوح
الخالصة
ا سن آم د بأستخدام قاعدة جورجن الي الموجب و الليكن ون األنتق سرت ف تم أقتراح نوع األرتباط بين األي
اح د بمعامل راآ ذرات المانحة في الليكن ز وال االت بأ طبيعة التآصر بين أيون الفل د األنتق م تحدي ستخدام و ت
ل ط أوريك ط مخط اطة مخط سابها بوس م أحت ي ت دي والت ال الليكن اييس المج ا(ومق وآانو-تناب ع ) س لجمي
. المعقدات المحضرة
بسم اهللا الرحمن الرحيم
.مريم علي سعيد البياتي: األسم
1984: المواليد
134 /د , 49 /ز , 836 / م- الدورة -بغداد : عنوان السكن
7713573: رقم الهاتف
[email protected] : األيميل
.ماجستير علوم آيممياء: التحصيل الدراسي
. جامعة النهرين– آلية العلوم – 2005بكالوريوس : سنة التخرج
.دآتور هادي محمد علي: أسم المشرف
15/05/2008: تأريخ المناقشة
تشخيص طيفي ألمالح بعض العناصر األنتقالية : عنوان األطروحة.في منصهرات األمالح بدرجة حرارة الغرفة
Curriculum Vitae
Name: Mariam Ali Saeed Al-Bayatie.
Year of Birth: 1984.
Address: Iraq- Baghdad, Al-Dura, 836/49/134.
Phone Number: 0096417713573.
E-mail: [email protected]
Academic Qualifications: Bs.c. in Chemistry (2005), Master degree in chemistry (2008), Al-Nahrian University, College of Science.
Supervisor: Dr. Hadi M.A. Abood.
Discussion Date: 15/5/2008.
Thesis title: Ultraviolet visible investigation of some transition metal compounds in room temperature ionic liquids, molten salts.