SOME TRANSITION METAL COMPOUNDS IN ROOM TEMPERATURE IONIC LIQUIDS MOLTEN...

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

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XVIII

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1

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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


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.

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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.


Nitrate.

3.2.2.3. Iron (III) chloride in Choline



3.2.3. Spectroscopic investigation of Cobalt (II) chloride

solution.

3.2.3.1. Cobalt (II) chloride in Choline


liquid solution.


Nitrate.

3.2.3.3. Cobalt (II) chloride in Choline


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VII


3.2.4. Spectroscopic investigation of Nickel (II) chloride

solution.

3.2.4.1. Nickel (II) chloride in Choline


liquid solution.


Nitrate.

3.2.4.3. Nickel (II) chloride in Choline



3.2.5. Spectroscopic investigation of Copper (II) chloride

solution.

3.2.5.1. Copper (II) chloride in Choline


liquid solution.


Nitrate.

3.2.5.3. Copper (II) chloride in Choline



3.3. Conclusion.

3.4. Suggestion for future work.

References.

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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

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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


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

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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.

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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


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.


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


temperature with addition of NH4NO3.

Figure (3-21):- The U.V.-Vis. spectra of CoCl2.6H2O of


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

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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



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.

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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.

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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


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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


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


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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

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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)


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


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


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).


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

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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

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[emim]

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6

h their go

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etc, are w24)

n of ionicre one or b

ese factors

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m chlorid

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nt. Examp

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).(24) For

num(III)

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roperties

and 1,3-

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for the

-ethyl-3-

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With

solvents

polar or

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crystalli

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mp=80o

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Fig.(1-2):- E

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low viscos

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with a lar

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ologies.(27)

al point of

thout solv

Note that

erefore mu

mic and en

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iquid com

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ompounds

xtures of

p=352o C)

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rge range

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Chapte

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1.6. A RTI

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1. Lo

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e liquids h

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aracterizue array o

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and capabi(33)

8

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o –81o C)

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iquids hav

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Introd

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in some

molecular

volatility

ake them

vents are

g or non-


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:


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


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


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

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however

the cation

me types of

based on thbased on thbased on th

zation wit

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Excluding

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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

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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.


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


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


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


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


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).


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


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)


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


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


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)


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-


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


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,


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


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


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


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


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)


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


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.


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.


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)


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).


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


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


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.


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


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) <


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


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/


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


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


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


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


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


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


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


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


80

3.2.3.3. Cobalt (II) chloride in Choline chloride/Malonic acid room


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


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


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

.


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


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


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


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].


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


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


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.


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/


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



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.


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.


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.

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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.

SOME TRANSITION METAL COMPOUNDS IN ROOM TEMPERATURE IONIC LIQUIDS MOLTEN...

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