Self Incompatible Solvent - monarch.qucosa.demonarch.qucosa.de/fileadmin/data/qucosa/documents/6042/data/diss.pdf · Self Incompatible Solvent von der Fakultät für Naturwissenschaften

Self Incompatible Solventvon der Fakultät für Naturwissenschaften der Technischen

Universität Chemnitz genehmigte Dissertation zur Erlangung des

akademischen Grades

doctor rerum naturalium

(Dr. rer. nat.)

vorgelegt von

M.Sc. Joanna M cfel-Marczewski

geboren am 06.08.1980 in Bydgoszcz (Polen)

eingereicht am 27.07.2009

Gutachter:

Prof. Dr. Werner A. GoedelProf. Dr. Stefan Spange

Tag der Verteidigung: 12.02.2010http://archiv.tu-chemnitz.de/pub/

http://archiv.tu-chemnitz.de/pub/

Bibliografische Beschreibung und Referat

ii

compound 3

"neutral"

compound 1 compound 2Self Incompatible Solvent

favorable interaction

unfavorable interaction

"neutral"

compound 3

"neutral"




"neutral"

compound 3

"neutral"




"neutral"

Bibliografische Beschreibung und Referat

Selbstinkompatibles Lösungsmittel

Joanna M cfel-MarczewskiTechnische Universität Chemnitz, Fakultät für Naturwissenschaften

Stichworte: Selbstinkompatibles Lösungsmittel, Mischungsthermodynamik, binäre & ternäreMischungen, Löslichkeitsparametern, Lösungskalorimetrie

In dieser Arbeit wird das neue Prinzip

der „Selbstinkompatiblen Lösungsmittel“

vorgestellt. Es wird theoretisch

abgeleitet, dass eine Mischung aus zwei

Substanzen mit ungünstigen

Wechselwirkungen bereitwillig eine

weitere Substanz aufnehmen sollte, die

diese ungünstigen Wechselwirkungen

durch Verdünnen vermindert. Dies sollte

umso stärker ausgeprägt sein, je

ungünstiger die Wechselwirkungen zwischen den beiden ersten Substanzen sind. Da sich

jedoch Substanzen mit sehr ungünstigen Wechselwirkungen physikalisch nicht mischen,

entstand die Idee, diese Substanzen durch eine kovalente Bindung aneinander zu binden. Ein

solches Molekül, das aus zwei inkompatiblen Hälften besteht, wird im Folgendem

Selbstinkompatibles Lösungsmittel genannt. Die in dieser Arbeit gewählten Substanzen

zeigen mäßige Inkompatibilität, deshalb ist ein Vergleich zwischen einfachen physikalischen

Mischungen und kovalent verknüpften Molekülhälften noch möglich. Dieses Prinzip wird für

binäre und ternäre Mischungen quantitativ berechnet und experimentell in drei Serien von

Experimenten bestätigt: i) unter Verwendung von Lösungskalorimetrie und Bestimmung der

Wechselwirkungsparameter zwischen Komponente 3 und einer bereits hergestellt

physikalischen binären Mischung aus Komponente 1 und 2, ii) unter Verwendung von

Lösungskalorimetrie und Bestimmung der Wechselwirkungsparameter zwischen Komponente

3 und den selbstinkompatiblen Losungsmitteln, die den in (i) gewählten Mischungen

entsprechen und iii) aus der Sättigungslöslichkeit der Komponente 3 in den entsprechenden

selbstinkompatiblen Lösungsmitteln. In diesen drei verschiedenen Messserien wird stets der

gleichen Trend beobachtet: Die Selbstinkompatibilität eines Lösungsmittels begünstigt den

Lösevorgang.

Abstract

iii

compound 3

"neutral"




"neutral"

compound 3

"neutral"




"neutral"

compound 3

"neutral"




"neutral"

Abstract

Self Incompatible Solvent

Joanna M cfel-MarczewskiChemnitz University of Technology, Faculty of Natural Science

Keywords: Self Incompatible Solvent, thermodynamic of mixing, binary & ternary mixtures, solubility parameters, Solution Calorimetry

In this thesis a new principle of Self

Incompatible Solvent is introduced. It is

shown theoretically that a preexisting

mixture of two substances (compound 1

and 2) with unfavorable interactions will

readily dissolve a third compound because

it diminishes the unfavorable interaction

between the compound 1 and 2 by dilution.

This behavior should be the stronger the

more unfavorable the interactions between

compound 1 and 2 are. However, substances with strong unfavorable interactions will not

mix. Therefore the idea pursued here is to enforce the desired preexisting mixture for example

by linking compound 1 covalently to compound 2. Such a molecule that is composed of two

incompatible parts is called Self Incompatible Solvent in this work. In this thesis examples of

incompatible compounds that show moderate incompatibility are chosen, therefore it was

possible to do a comparison between simple physical mixtures and covalently linked

incompatible molecules. The theoretical prediction of the theory is compared with

experiments. This principle is calculated quantitatively for binary and ternary mixtures and

compared with the experimental results in three distinct series of experiments: i) by using

solution calorimetry and calculation of the interaction parameters between compounds 3 and

the preexisting binary mixture of compound 1 and 2, ii) by using solution calorimetry and

calculation of the interaction parameters between compound 3 and the Self Incompatible

Solvent that correspond to the mixtures used in (i) and iii) from the saturation solubility of

compound 3 in the Self Incompatible Solvent. The results obtained from the theoretical

prediction and these obtained from the three different series of experiments show the same

trend: the self incompatibility of the solvent improves the dissolution process.

Table of contents

iv

Table of contents

Bibliografische Beschreibung und Referat ....................................................................... ii

Abstract................................................................................................................................... iii

List of abbreviations............................................................................................................ vii

Chapter 1: Thermodynamics of mixing............................................................................ 1

1.1 Introduction .................................................................................................................... 2

1.2 Entropy of mixing ........................................................................................................... 7

1.3 Enthalpy of mixing ......................................................................................................... 9

1.4 Gibbs free energy.......................................................................................................... 10

1.5 Group contribution methods.......................................................................................... 12

Chapter 2: Self Incompatible Solvent .............................................................................. 16

2.1 The principle................................................................................................................. 17

2.2 The theoretical predictions ............................................................................................ 19

2.3 The strategies for experimental verification................................................................... 24

Chapter 3: Model systems .................................................................................................. 26

3.1 Model substances.......................................................................................................... 27

3.1.1 Active ingredients: Compound 3..................................................................... 27

3.1.2 Model substances: Compounds 1 and 2........................................................... 28

3.1.3 Strategies of getting results ............................................................................. 30

3.2 Physical 50/50 mixture.................................................................................................. 33

3.2.1 Results............................................................................................................ 34

3.3 Self Incompatible Solvent ............................................................................................. 37

3.3.1 Results............................................................................................................ 38

3.4 Conclusion and outlook................................................................................................. 40

Chapter 4: Mixing thermodynamics of binary mixtures .............................................. 42

4.1 Compound 1 in Compound 2 ........................................................................................ 43

4.2 Dilution experiment and Hess law................................................................................. 44

4.3 Results for binary mixtures ........................................................................................... 46

4.4 Results using various solvents....................................................................................... 49

Table of contents

v

4.5 Comparison of the experimental data with the existing theories..................................... 51

Chapter 5: Saturation solubility ........................................................................................ 53

5.1 Saturation solubility ...................................................................................................... 54

5.2 Extracting ki3 from saturation solubility ........................................................................ 55

5.2.1 Ideal solution.................................................................................................. 55

5.2.2 Regular solution ............................................................................................. 56

5.3 Solubility experiment .................................................................................................... 58

5.4 Results .......................................................................................................................... 59

5.4.1 Precision solution calorimeter versus Saturation solubility.............................. 59

5.5 Conclusion.................................................................................................................... 62

Chapter 6: Polymers ............................................................................................................ 63

6.1 Motivation .................................................................................................................... 64

6.2 Polybutadiene-block-poly(2-vinylpyridine) ................................................................... 65

Chapter 7: Experimental part............................................................................................. 71

7.1 Purification process of active ingredients ...................................................................... 72

7.1.1 Anthracene .................................................................................................... 72

7.1.2 9-Anthracene carboxylic acid ......................................................................... 73

7.1.3 Acridine ......................................................................................................... 73

7.2 Used equipment ............................................................................................................ 74

7.2.1 Nuclear Magnetic Resonance (NMR) ............................................................. 74

7.2.2 Infrared Spectroscopy (IR) ............................................................................. 74

7.2.3 Fluorescence Spectroscopy............................................................................. 74

7.2.4 Precision Solution Calorimetry....................................................................... 76

7.2.4.1 The principle of calorimetry.................................................................. 76

7.2.4.2 SolCal Precision Solution Calorimeter .................................................. 77

7.3 Synthesis and purification of the model substances for compounds 1 and 2 ................... 81

7.4 Synthesis and purification of the model substances for Self Incompatible Solvents ....... 91

7.5 Results from measurements with the SolCal Calorimeter ........................................... 109

7.6 Polymers..................................................................................................................... 118

Appendix 1 ....................................................................................................................... 126

Appendix 2 ....................................................................................................................... 128

Table of contents

vi

Chapter 8: Conclusion ...................................................................................................... 130

References ........................................................................................................................... 133

Acknowledgments ............................................................................................................. 136

Curriculum Vitae ............................................................................................................... 137

Selbstständigkeitserklärung ............................................................................................. 139

List of abbreviations

vii


Abbreviation ExplanationFirst

mentionedon page:

a Activity coefficient 57 c Concentration 76

12c Constant describing the mutual interaction of twosubstances

19

DP Degree of polymerisation 67cohE Cohesive energy 6

dE Cohesive energy of contribution of dispersion forces 13

hE Cohesive energy of contribution of hydrogen bonding 13

hiE Hydrogen bonding contribution 13

pE Cohesive energy of contribution of polar forces 14.eqn Equation 3

diF Molar attraction of dispersion group 13

piF Molar attraction of polar group 13

tF Molar attraction function 13

pF Polar components of molar attraction function 13

Gmix Gibbs free energy of mixing 3Gsol Gibbs free energy of solubility 10Hmix Enthalpy of mixing (two liquid substances) 3Hsol Enthalpy of solubility (one compound crystalline and

one compound liquid)11

Hvap Enthalpy of vaporization 6

3HfusFusion enthalpy, enthalpy of melting of compound 3,heat of fusion

11

IR Infrared 74ijk Equal to k12, k13, k23, ksis3

2ik Interaction parameter between compound “i” andcompound 2

11

3ik Interaction parameter between compound “i” (i is aplaceholder for either 1, 2 or SIS) and compound 3

11

12k Interaction parameter between compound 1 andcompound 2

19


21


21

3sisk Interaction parameter between the Self IncompatibleSolvent and compound 3

24

3-mixture50/50k Interaction parameter between 50/50 mixture ofcompound 1 and 2 and compound 3

33

Bk Boltzman constant 7


viii

m Multiplied 85m Weight 76Mw Molecular mass 67n Number (e.g. of molecules) 8N Molar amount 8

iN Molar amount of compound “i” 13

3N Molar amount of compound 3 56NMR Nuclear Magnetic Resonance 74p Pressure 12

P Power 80PDI Polydispersity Index 67R Gas constant 6s Singulet 82

Smix Entropy of mixing 3

3SfusFusion entropy, entropy of melting of compound 3 11

Ssol Entropy of solubility (one compound crystalline and onecompound liquid)

11

SIS Self Incompatible Solvent 17 t Triplet 85t Time 80T Temperature [K] 3

bT Boiling point 14

crT Critical temperature 14

T Lydersen constant of liquids 14Uvap Energy of vaporization 12

UV Ultra violet 6V Volume 6

V Change of volume 12*V Volume of a given binary mixture (or the Self

Incompatible Solvent)20

V Molar volume 55

iV Molar volume of compound “i” (i is a placeholder foreither 1, 2 or SIS)

57

3V Molar volume of compound 3 57x Molar fraction 57

Polarizability 5Volume fraction 6

*1

Volume fraction of compound 1 in a given binarymixture (or part 1 in the self incompatible solvent)

20

*2

Volume fraction of compound 2 in a given binarymixture (or part 2 in the self incompatible solvent)

20

Chemical shift 74Solubility parameters 6

D Solubility parameters for dispersion forces 6P Solubility parameters for polar forces 6H Solubility parameters for hydrogen bonding 7


ix

Standard potential 57

3sol Change in chemical potential of compound 3 upondissolution

56

)(3 s¤ Chemical potential of the undissolved pure solidcompound 3

54

)(3 l¤ Chemical potential of the dissolved compound 3 54Number of possibilities to realize a given state of asystem

7

Parameter from Flory – Huggins theory 20

Chapter 1Thermodynamics of mixing

1

Chapter 1:Thermodynamics of mixing

Abstract

This chapter summarizes the basic information about the interaction between molecules and

the thermodynamics of mixing. Therefore, such terms like the entropy of mixing, the enthalpy

of mixing or Gibbs free energy will be defined. Additionally two group contribution methods,

namely Hoftyzer - van Krevelen and Hoy will be described in details and finally the solubility

parameters calculated on the examples of molecules via these two methods will be compared.


2

1.1 Introduction

In our daily life liquid preparations are universally used and very favorable. For

example we usually use liquid soap instead of solid soap, toothpaste instead of toothpowder or

mineral oil instead of coal. The most common liquid preparations are dispersions: something

in form of small particles “embedded” in a liquid phase. Dispersions have found wide range

of different applications in almost all foodstuffs (milk, yoghurt, ice-cream),[1][2] but also in

paints[3], personal care products or pharmaceutical[4] and agricultural formulations[5]. The

main advantages of dispersions are: a high “solid content” at a low viscosity and that they are

environmentally friendly, because in most cases water can be used as a liquid component.

However, dispersions have also some disadvantages like light scattering and that they are

metastable and in their metastable states they can lead to creaming, Ostwald ripening or

coagulation[6]. Another kind of liquid preparations are solutions: a mixture of several

compounds that is homogeneous down to the molecular level. A simple, constructive

experiment helps to clarify the differences between a solution, dispersion and not soluble

substances. Three glass containers are partially filled with warm water which should act as the

solvent in case of solutions or as a liquid phase in case of dispersion. Into the first container

sand is added, into the second one clay and into the third one table salt. All of these

compounds are added in the same amount. After the addition three different behaviors are

observed. In the first container, the sand is not dissolved and it settles down at the bottom.

The clay in the second container does not settle immediately but forms in this case a

metastable dispersion that needs a day to completely settle down at the bottom of the

container. In the third container, the salt is completely dissolved in water, forms solution,

which does not phase separate regardless of how long we wait, see Figure 1.

sand silt saltsandsand siltsilt salt

Figure 1: Schematic illustration of three different behaviors in a liquid phase.

sand clay table salt


3

Exactly dispersions and solutions are the reason why our live become easier and more

comfortable, they help us to clean our homes, make paint flow, ink dry or keep bridges from

rusting. Without liquid preparations it will not be possible to use some of products or their

performance will not be so good.

As mentioned before, dispersions are usually metastable. Thus, if one has to optimize

their stability, one usually has to influence the kinetics of the destructive process. Solutions in

contrast are in general thermodynamically stable. Thus, it is worth to give their equilibrium

thermodynamics a close look. The Gibbs free energy ( Gmix ) is the driving force of the

dissolution processes and the value of Gmix must be below zero for spontaneous mixing[7].

The Gibbs free energy depends on the enthalpy of mixing ( Hmix ) and the entropy of

mixing ( Smix ), see eqn. 1

STHG mixmixmix eqn. 1

Thus, Gmix becomes negative if Smix is positive and Hmix is negative or at least if one of

the two terms on the right hand side is negative and dominates the other one. It is easy to

justify that very often Smix is positive. If two initially pure substances are mixed the number

of possible arrangements of the molecules in space tremendously increases and thus the

entropy of the system. Hmix , on the other hand, very often is positive, it might become

negligible small but still remains positive if two similar substances are mixed. Two very

similar substances have a tendency to mix, because of the small value of Hmix . This

relationship was already discovered by the alchemists, who already framed the famous rule

“similia similibus solvuntur” which means “like dissolves like”.

One might ask why there are only a few examples of negative mixing enthalpies. To

answer that question, let us have a close look at the interactions between individual

molecules. First there are so called “short range interactions”. Interactions, that may be

assigned to overlap of quantummechanical wave functions, but can not be considered as

covalent bonds. Examples are: (i) hydrogen bonds, which are attractive interactions that occur

when hydrogen (H) is covalently connected to highly electronegative elements (A) and

interacts with another highly electronegative atom (B) like oxygen, nitrogen or fluor which

have a free pair of electrons (A-H......B), (ii) donor-acceptor interactions, where one molecule

is providing an occupied orbital that overlaps with an unoccupied orbital of the acceptor (iii)

charge – transfer interactions, where there is an energetically favorable resonance hybrid


4

between two states that differ from each other by the transfer of electrons from one molecule

to the other. In the case of short range interactions the enthalpy of the interaction depends on

the mixing of wave functions and is often most favorable if two wave functions of different

characters interact (e. g. an occupied orbital with an unoccupied orbital, a -bond to hydrogen

with an occupied non binding orbital). In general if the enthalpy of the interaction of a pair of

molecules is a function of the one property of the first molecule (e.g. its acidity) and a

property of different character of the second molecule (e.g. its basicity) we call this

interaction non symmetric. Non symmetric interactions can give rise to negative enthalpies of

interaction, as anyone can confirm who ever has mixed sulfuric acid with water.

Besides the quantummechanical short range interactions, there are other interactions

that can be described qualitatively by classical physics and more precisely by Coulomb

interactions[8]. First of all, there are the Coulomb interactions between ions or charged

molecules. These interactions can lead to negative enthalpies if the ions have opposite charge.

Besides these, there are as well coulombic type interactions if the molecules are not charged,

but have a dipole moment. Interactions in which a favorable alignment of dipoles give rise to

a favorable coulombic interaction are summarized as dispersion interactions or van der Waals

interactions[9][10]. In detail, one can distinguish between three types of dispersion interactions:

(i)London dispersion forces (induced dipole – induced dipole)[11]. These forces occur when

two molecules are in proximity and the random fluctuations of their dipole moments become

correlated. These two molecules attract each other, that means that the nucleus of one

molecules interacts with the electron of the other one and vice versa. This causes a creation of

temporary dipoles. The intermolecular attractions are greater between large molecules

because the number of temporary dipoles is also greater. Although the London forces are

present in all molecules, they do not play the dominant role in polar molecules. Polar

molecules orient themselves preferentially with antiparallel orientation of their dipole

moments and this causes further increase of the intermolecular attraction. If both considered

molecules have a permanent dipole moment, these symmetrical interactions are called (ii)

Keesom interactions (dipole-dipole) and depend on the temperature. With increasing

temperature, the rotation of the molecules also increases and this causes a decrease of Keesom

interactions. The third kind of interactions (iii) Debye forces (dipole – induced dipole) occurs

for example when a nonpolar molecule is in proximity with polar molecules. This causes an

induced polarisation of the nonpolar molecule by the nearby permanent dipole of the other

molecule. Such induced dipole is not affected, when the temperature is increased and thus the

rotation of the molecules is more pronounced. That means that the Debye interactions do not


5

depend on the temperature as strong as the Keesom interactions. Both, the London and the

Keesom interactions are symmetric. We call interactions symmetric, if the enthalpy of the

interaction of a pair of molecules depends on a certain property of the first molecule and

exactly the same property of the second molecule, for example to the strength of the dipole

moments of both molecules. Very often the strength of symmetric interactions is given by the

product of the values of this property of the first molecule and the corresponding value of the

second molecule. Such interactions usually give rise to negative enthalpies of pairwise

interactions. However, the enthalpy of mixing - viz. the difference of the enthalpy of

interaction between various molecules in the mixture minus the enthalpy of interaction of the

molecules in the pure substances before mixing – is positive. The absolute value of the

enthalpy of mixing will be the smaller, the more similar the regarded molecules are. If, for

example, equal volumes of two substances are mixed together the interactions of these

substances in pure state and as mixture as a function of their polarisabilities, 1 and 2 are

shown by the equations below, which are simple numerical examples:

Interaction enthalpy in the mixture: 212

Interaction enthalpy in pure substances: 22

21

Enthalpy of mixing = Interaction in the mixture - Interaction in pure substances:

02 221

22

2121

The energy of interaction of these molecules among each other is always negative (the

interaction is favorable). However, the difference of the energy of interaction in mixture and

in pure substances is always positive or zero, that means unfavorable for the mixing of the

substances. These unfavorable interactions are smaller when the substances are similar to each

other. Therefore the rule applied by alchemist already mentioned “similia similibus

solvuntur”, seems to work, because polar substances are good soluble in polar solvents and

unpolar substances are good soluble in unpolar solvents. Due to this fact, it can be deduced

that the strength, density, mobility and induction of dipoles give information about the

enthalpy of mixing.

In the thermodynamics of mixing, all above mentioned interactions need to be

considered. Although the short range interactions often give rise to comparatively high

interactions energies if only a single pair of molecules is considered, the full interaction

energy in a bulk medium is obtained from integration over the whole volume. It is in the

nature of the long range interaction, that they may be weak but extended over comparatively

long distances. Thus, in most cases long range interactions are the dominating term if the

whole volume is considered.


6

The above mentioned properties of dipole moments of molecules are connected with

the term “polarity”. This parameter is elusive but very important and therefore a lot of polarity

scales were developed. One of such scales is the relative permittivity (dielectric constant). The

substance is classified as non polar when the dielectric constant is less than 15[12] and is not

miscible with water. For example, hexane belongs to this group with a dielectric constant of

1.88[13] or tetrahydrofurane with a dielectric constant of 7.4[14]. Examples for polar solvents

are water with a dielectric constant of 78.39[15] or methanol with a dielectric constant of

32.70.[16] Another way to measure the polarity of solvents is the dipole moment. Molecules

with a large dipole moment and a high dielectric constant are classified as polar. Molecules

with a low dipole moment and a small dielectric constant are classified as non polar. One of

the polarity scales is the so called donor-acceptor. It is applied when a solvent interacts with

some characteristic substances like for example strong Lewis bases or strong Lewis acids.

Another one is solvatochromism. In that case is the wavelength of the UV absorption maxima

of a suitable dye (usually the excited state has a dipole moment significantly larger or smaller

than the ground state) experiences a significant shift if is dissolved in solvents with various

polarities.

On the other hand quite successful theories have been developed to describe the

enthalpy of mixing directly, deducing the pairwise interactions from suitable constants that

characterize individual compounds (not pairs) but can not necessary termed “polarity”. The

most prominent systems of this kind are the Hildebrand and the Hansen system.[17][18] The

Hildebrand parameter of one substance is equal to the square root of the energy of evaporation

(called as well cohesive energy)* divided by the specific volume:

21

VEcoh eqn. 2

The enthalpy of mixing of two substances is equal to a term which depends on the two

volume fractions of the considered substances multiplied by the square of the difference of the

Hildebrand parameters:

212

21)21( /VHmix eqn. 3

One advantages of the Hildebrand parameter is that it can be directly measured for

vaporizable substances. However, this treatment reduces the interactions between molecules

to one kind of interactions, although there are several ones like Keesom, Debey or London

* RTHE vapcoh


7

interactions. Therefore Hildebrand’s concept was extended by Hansen. Hansen split the

solubility parameter up in three types which he called: “dispersion forces ( D )”, “polar forces

( P )” and “hydrogen bonding ( H )”. The Hildebrand parameter is expected to be given by

Hansen’s parameters according to:2222 HPD

eqn. 4

The enthalpy of mixing in the Hansen system is described as following:

212

212

212

21)21( / HHPPDDmix VH eqn. 5

It is worth noting that Hansen’s distinction between “dispersion” and “polar” forces is not

identical to the often used definition of dispersion forces, which comprises polar interactions

within the dispersion forces. If similarities to the definition given in page 5 are sought,

Hansen’s dispersion parameter has most similarity to London interactions and his

“dispersion” interactions seem to match with Debey and Keesom interactions. It is further

worth noting that Hansen assumes that hydrogen bonding can be described as symmetric

interactions although it is not symmetric.

The Hildebrand and Hansen systems are not perfectly matching observations but are

quite efficient and thus very often used. Accordingly, there is a lot of effort to find solvents

which are similar to the solute, because in such case a minimization of the unfavorable

interactions is achieved. The alchemist’s sentence seems to work if the symmetric interactions

are dominating, but still question arises: Is it possible to get a negative enthalpy of mixing

( OHmix ) even if symmetric interactions occur? The goal of this thesis is to show that,

yes, this may be possible, namely if the solute is dissolved not in one solvent but in a

preexisting mixture of solvents which “do not like each other”.

1.2 Entropy of mixing

As it was mentioned in section 1.1, a mixing of two substances causes changes in

enthalpy and entropy of the system. These both terms are responsible for changes of Gibbs

free enthalpy (see eqn.1, section 1.1) and the system leads to a decrease of the value of this

term. Following the Boltzman expression, the entropy of a system is proportional to the

logarithm of the number of various possibilities of states of the system which can be realized,

so called microstates ( ), see eqn. 6 and eqn. 7.


8

lnBkS eqn. 6

lnstart

endBkS

eqn. 7

In the easiest case n molecules of an ideal gas are distributed in the volume of V lattice site. In

case of ideal gas double occupancy is allowed and each of the molecules has V lattice sites

available and is equal to V in the power of n (eqn. 8).

~ nV eqn. 8

When two such gases are mixed, n1 molecules in volume V1 are expanded into the additional

volume V2 and at the same time n2 molecules in volume V2 are expanded into the additional

volume V1. Then each kind of molecules has a higher volume which can be occupied. The

changes of entropy caused by this process are proportional to the negative logarithm of the

volume fractions of each gas after mixing (eqn. 9, eqn. 10, eqn. 11).

111

11

1

1

211 ln1lnln RNknn

V

nVVkS BBmix

eqn. 9

2211 lnln RNRNSmix eqn. 10

0lnln 22

21

1

1 RV

RVV

Smix with 121 eqn. 11

In this work it is dealed with liquids and not with gases and therefore double occupancy of

lattice sites is not allowed. This forbiddance of double occupancy is related to the pure

substances and to the mixtures. In such cases the eqn. 8 becomes more complicated, see

eqn. 12.

1........21 nVVVV eqn. 12

These complications arising from the impossibility of double occupancy, however, have

identical effect on the entropy of the initial stage as well as on the entropy of the final stage.

In this case only the change of entropy upon mixing is for interested, thus these mathematical

terms that would be needed to describe these complications can be canceled and thus, the final

results of Smix is identical to the result describing the mixing of ideal gases (eqn. 11).

Graphic illustration of the entropy of mixing is shown in Figure 2. The entropy of mixing is


9

always positive and has a maximum in 50/50 mixtures if the molecular volume of both

compounds is identical.

0 0.2 0.4 0.6 0.81

0

1

2

)

10

VSsol

0 0.2 0.4 0.6 0.81

0

1

2

)

10

VSsol

0 0.2 0.4 0.6 0.81

0

1

2

)

10

VSsol

0 0.2 0.4 0.6 0.81

0

1

2

)

10

VSsol

Figure 2: The entropy of mixing of a binary mixture as a function of the volume fraction of one of its compounds.

1.3 Enthalpy of mixingThe enthalpy of mixing gives information about the interaction between two

molecules. The interactions can be favorable then the value of the enthalpy of mixing is

negative ( 0Hmix , exothermic). When the interactions are neutral the value of enthalpy of

mixing is near zero ( 0Hmix , athermal). The interactions can also be unfavorable and the

value of the enthalpy becomes positive ( 0Hmix , endothermic). In general the enthalpy of

mixing is proportional to the product of the volume fractions of the two compounds that are

mixed. An illustration of the enthalpy of mixing of binary mixtures is shown in Figure 3. In

general, almost all solubility processes are endothermic. This can be quantitatively described

by the mean field theory if it is assumed that all interactions between the molecules are

symmetric. The mean field theory is described and used for theoretical prediction in the next

chapter.


10

0 0.2 0.4 0.6 0.81

0

1

2

)

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VHsol

Figure 3: The enthalpy of mixing of a binary mixture as a function of the volume fractionof one of its compounds.

1.4 Gibbs free energyThe combination of enthalpy and entropy of mixing is given by the Gibbs free energy. If the

entropy of mixing is subtracted from the enthalpy of mixing then the graphic illustration of

Gibbs free is shown in Figure 4

0 0.2 0.4 0.6 0.81

0

1

2

10

VGsol

miscibility gap

0 0.2 0.4 0.6 0.81

0

1

2

10

VGsol

miscibility gap

Figure 4: The Gibbs free energy of mixing.

The mathematical description of the Gibbs free energy of mixing is the following:


11

2222

2 lnln iiii

isol kRTV

RTVV

G

In Figure 4 it can be seen that the system is separated into two phases, if there is a straight line

that has two points in common with the curved line but is situated below it. The lowest of

these lines is usually given by a line that is a tangent to the curve at two non-identical points.

These points indicate the composition of the coexisting phases.

Till now the thermodynamics of mixing were described for the case of two liquids. If

one of the mixed compounds is crystalline, the eqn. 13 for the Gibbs free energy of mixing

must be modified and the product of the heat of melting of this crystalline substance and its

volume fraction must be added. In this case the equation for the Gibbs free energy of mixing

is given in the eqn. 14:

The corresponding diagrams for the entropy, the enthalpy and the Gibbs free energy of mixing

if one of the mixing compounds is crystalline and the another one is liquid are shown in

Figure 5.

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

a.) b.) c.)

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

a.) b.) c.)a) b) c)

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

a.) b.) c.)

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

a.) b.) c.)a) b) c)

Figure 5: Graphic illustrations of a) entropy b) enthalpy and c) Gibbs free energy of mixingif one of the mixed compounds is crystalline.

* 22

22

222

Hi

HPi

PDi

Dik

* eqn. 13

33

3

3333

3

33

3 lnlnV

Hk

VS

RTV

RTVV

G fusii

fusi

i

isoleqn. 14

HSTG

HSTG


12

1.5 Group contribution methods

In section 1.1 it was already mentioned that there are two most common theories: the

Hildebrand and the Hansen, which described solubility behavior via solubility parameters. In

Hildebrand and Hansen theories the solubility parameter ( ) is correlated directly to the

cohesive properties of the solvents and it is calculated from the cohesive energy:

RTHVpHUE vapvapvapcoh eqn. 15

Nevertheless, the parameter proposed by Hildebrand is of good predictability only for non

polar solvents with low molecular masses. His concept was successfully extended by Hansen,

who proposed three parameters to described solubility behavior as a sum of what he called:

dispersion forces, polar forces and hydrogen bonds (see eqn. 5, section 1.1). Hansen

parameters can experimentally be obtained for existing compounds by means of solubility or

swelling or calorimetry, using a suitable number of pairwise correlations. Although, these

two valuable methods have found widespread applications, especially for solvent selections in

polymeric systems, it is worth to mentioning, that these methods also have a limitation, for

example because the solubility parameters are not tabulated for all substances and can not be

obtained for unknown substances. Therefore, it was important to develop another method

which can estimate and predict solubility behavior from molecular structures. One such

method is the so called “group contribution method” and is based on the assumption that

distinct functional groups always contribute in the same manner to the interaction parameters,

even if they are parts of various molecules. Each functional group is assigned to an individual

parameter, the group contribution. The interaction parameter of a complete molecule is

obtained by 'summing up' the individual parameters of all groups comprising this molecule.

The big advantage of this method is its simplicity, 'summing up' the group contribution allows

to predict solubility. If these methods are precise, they can replace measurements or they

make the comparison of the experimental data with the theoretical calculations possible.

Usually these group contributions are derived from experimentally obtained solubility

parameters of a significant number of molecules via multi-parameter correlations that result in

sets of parameters that give minimized deviation between prediction and experiment. There

are various group contribution methods, which differ from each other a) by the way to classify

functional groups and b) by the mathematical equations used to calculate solubility

parameters from the group contributions. As it can be seen from the summaries of some of

these methods given below, the mathematical procedures may look not so simple and the

justifications for the used dependencies not always seem to be based on ab initio principles,


13

but may as well be to some extend empiric. For example one group contribution method was

proposed by Hoftyzer and van Krevelen. It is applied to calculate the dispersion, polar and

hydrogen bonding of compounds by equations given below:

VFdi

d

eqn. 16

V

Fpip

2 eqn. 17

VEhi

h

eqn. 18

Where F is the molar attraction constant introduced by Small, correlated to the cohesive

energy, eqn. 19:

21

298VEF coh

eqn. 19

and diF , piF , hiE are the dispersion group molar attraction, the polar group molar attraction

and the hydrogen bonding contribution, respectively. As can be seen, this method is based on

the eqn. 5 of Hansen, which means that the cohesive energy is a sum of the contribution of

dispersion forces ( dE ), the contribution of polar forces ( pE ) and the contribution of hydrogen

bonding ( hE ):

hpdcoh EEEE eqn. 20

The final value of the predicted solubility parameter is calculated from equation given below:

222hpd

eqn. 21

Besides this method of Hoyzer and van Krevelen there are other methods known:

Fedors, Dunkel, Hayes or Hoy, which differ in the calculation procedure of the final solubility

parameter.

Hoy for example used a system of equations for the estimation of the solubility

parameters and its compounds. He proposed different equations for low-molecular liquids

(solvents) and for amorphous polymers. From the following equations the solubility

parameter for solvents by Hoy method can be calculated:

itit FNF ,eqn. 22

ipip FNF ,eqn. 23


14

iiVNV eqn. 24

iTiT N ,eqn. 25

tF , pF V , T are the molar attraction function, the polar components of molar attraction

function, the molar volume and the Lyderson correlation for non-ideality, respectively.

Hoy proposed the supporting equations given below:

VTTLog

cr

b log1585.039.3eqn. 26

2567.0 TTcr

b

TT eqn. 27

with bT as boiling point and crT as the critical temperature. Finally the -components can be

obtained and an overall value of the solubility parameter can be calculated, see the equations

below:

VBFt

t where 277Beqn. 28

21

1BF

F

t

ptp

eqn. 29

21

1th

eqn. 30

21

222hptd

eqn. 31

As can be seen, Hoy, proposed slightly complicated equations for the prediction of solubility

parameters in comparison to Hoftyzer and van Krevelen.

Nevertheless, the both calculations are expected to result in satisfying predictions by

their authors, and therefore should reasonably agree with each other. Therefore it is worth to

compare these two methods. For this purpose, a few molecules are chosen (see Table 3,

section 3.1.2, model substances for compound 1 and 2) and the solubility parameters of these

molecules are calculated by these two methods. The results are shown in Figure 6 .


15

0 2000 4000 6000 8000 10000

0

200

400

600

800

1000

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y H

ofty

zer -

van

Kre

vele

n m

etho

d

D1- D

2 )2+( P1-

P2)2+( H

1- H2)2 (kJ/m3)

calculated by Hoy method

0 50 100 150 200 250 300 3500

50

100

150

200

250

300

350

D 1-D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2(k

J/m

3 )ca

lcul

ated

by

Hof

tyze

r - v

an K

reve

len

met

hod

D1- D

2)2+( P1-

P2)2+( H

1- H2)2 (kJ/m3)


0 2000 4000 6000 8000 10000

0

200

400

600

800

1000

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y H

ofty

zer -

van

Kre

vele

n m

etho

d

D1- D

2 )2+( P1-

P2)2+( H

1- H2)2 (kJ/m3)


0 50 100 150 200 250 300 3500

50

100

150

200

250

300

350

D 1-D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2(k

J/m

3 )ca

lcul

ated

by

Hof

tyze

r - v

an K

reve

len

met

hod

D1- D

2)2+( P1-

P2)2+( H

1- H2)2 (kJ/m3)


Figure 6: Comparison of predicted solubility parameters calculated by the Hoftyzer-vanKrevelen method and the Hoy method[17].

In Figure 6 can be seen that the both calculations, which were developed independent of each

other make similar predictions, nevertheless with obvious limitations. The highest deviation

can be seen by molecules which contain fluor (Data points labeled , and , see Table 3,

section 3.1.2). In section 4.5 of this thesis the comparison of the group contribution method of

Hoftyzer-van Krevelen and Hoy with experimental results will be shown.

765

Chapter 2Self Incompatible Solvent

16

Chapter 2:Self Incompatible Solvent

Abstract

The mean field model of solution thermodynamics is described in chapter 1. If symmetric

interactions prevail, this model predicts that the enthalpy of mixing in binary systems is

always positive, at best can be zero. Substances chemically similar to the solvent are very

good soluble in it. In this chapter a new theory called Self Incompatible Solvents is described.

In general, this theory assumes that, if a solvent is composed of two distinct mutually

incompatible parts and an additional substance is dissolved in this solvent, that there is a

negative contribution to the enthalpy of mixing; which is the more pronounced the more

incompatible the two parts are. Finally a procedure to test these predictions experimentally is

developed.


17

2.1 The principle

In the last years some solvents showed surprising good dissolution properties. These

solvents are comparable in their chemical structure to detergents because they have

dissimilarities in the molecular structure. However, they can not be termed detergents because

the size of the molecule is much smaller than the molecule of detergents, thus the formation of

micelles at even low concentrations, which is typical for detergents, does not occur. These

molecules dissolve even polymers which are only partially or even not soluble in other

solvents. This suggests that the rule of thumb “like dissolved like” is not the only one rule

existing in solvent behavior. In Table 1 are shown examples of three groups of solvents.

Table 1: Groups of solvents.Simple solvents Solvents with some kind

of internalincompatibility

Amphiphiles(soaps)

Butoxyethanol

Methanol

Hexafluoroisopropanol

Sodium dodecylsulfate

In the group of simple solvents, methanol is used as an example, which is a common polar

protic solvent. Such solvents will dissolve other polar substances, which is in agreement with

the rule “like dissolves like”. To the amphiphiles (soaps) belongs for examples sodium

dodecylsulfate. These molecules consist of a long alkyl chain (unpolar, hydrophobic part) and

sodium sulfate group (polar, hydrophilic part). Such substances are called surfactants –

“surface active agents”, which often do not dissolve substances but form micelles and

“solubilize” substances within the micelle core. The middle column in Table 1 contains

solvents that comprise some kind of internal incompatibility, but can not be considered as

amphiphiles. For example butoxyethanol which consists of an unpolar part – left side - and a

polar part – right side of the molecule. Another example of this group is

hexafluoroisopropanol which is an acidic alcohol with strong hydrogen bonding and is known

as a good solvent in many polymer systems. This kind of solvents often seems to be “magic”.

It may dissolve substances including those that are not soluble in the most common organic

solvents. From now let us call solvents like these Self Incompatible Solvents (SIS).

OOH

F3C CF3

OH

H3C OH O S

O

O

O NaNa


18

In the following it is set out to explain the principle of self incompatible solvents, first

using a simple analogy and subsequently in quantitative thermodynamic terms. If two

substances with unfavorable interaction: like compound 1 and 2 (represented by two men in

Figure 7) are mixed together, one needs to overcome the positive heat of mixing. The addition

of a third compound (represented by a pretty young lady in Figure 7) into this preexisting

mixture decreases the unfavorable interaction between the compound 1 and 2. Thus, the

addition of a third compound to an existing binary mixture to create a ternary mixture may be

exothermic, provided that the interaction between the two initially mixed compounds is

strongly endothermic.

compound 3

"neutral"




"neutral"

compound 3

"neutral"




"neutral"

Figure 7: Schematic illustration of the principle of Self Incompatible Solvent.

On the other hand, if the interaction between compound 1 and 2 is as desired endothermic

these two incompatible compounds actually will not mix. However, this unfavorable mixture

can be enforced, e.g. by linking compound 1 covalently to compound 2 (symbolized by the

bench in Figure 7), as it actually is the case in the examples mentioned above. Such solvents,

that are composed of two incompatible parts that can not demix into a macroscopic phase due

to covalent bonds and will neither micro phase separate into e.g. lamellar morphology, are

called Self Incompatible Solvents. The goal of this dissertation is to go beyond this instructive


19

but purely qualitative explanation and to find a quantitative description, and correlate it to

experimental data.

2.2 The theoretical predictions

The thermodynamics of mixing can be described by a mean field model. As it is usual for

mean field models, the environment surrounding each molecule is assumed to be structure

less and its properties depends in a predictable way on the mean composition of the mixture.

In other words, it is assumed that the mixture of two compounds is homogeneous, no phase

separation occurs and the interactions of one molecule with its environment are a linear

function of the volume fractions of each of the compounds. If the interaction energy of unit

volume of substance 1 surrounded by pure substance 2 is proportional to a characteristic

constant 12c , the change in interaction energy upon mixing a volume V1 of substance 1 with a

volume V2 of substance 2 (interaction energy after mixing – before mixing) is given by:

mixingbeforemixingaftermix HHH 1221 eqn. 32

222111121222212211111221 cVcVccVccVHmixeqn. 33

Constant 12c describes the interaction of two substances with each other. The enthalpy of

mixing is equal to negative interactions of the pure compound 1 with itself before mixing,

plus interactions of compound 1 with itself and with compound 2 in the mixture, minus

interaction of pure compound 2 with itself before mixing, plus interactions of compound 2

with its self and with compound 1 in the mixture. The eqn. 33 can be simplified as given

below:

12122222222122111111111221 cccccc

VHmix eqn. 34

12122222122111111221 11 cccc

VHmix eqn. 35

cccV

Hmix222112211121

1221 2eqn. 36

221211211221 2 ccc

VHmix eqn. 37


20

12211221 k

VHmix with 22121112 2 ccck

eqn. 38

In the community dealing with low molar mass substances, the enthalpies of mixing often are

measured in this parameter k12, which has the dimensions of energy per volume. However, in

the polymer community it is common to use instead of the ratio of k12 to kBT, so called [19][20]

parameter, for the same purpose.

122112211221 Tkk

VH

Bmix eqn. 39

The next step is a qualitative description of the mixing phenomena based on the mean

field theory. As already discussed in chapter 1, mean field theory can provide predictions for

the energy of mixing and the entropy of mixing. It is assumed that the entropy of mixing is

favorable but not significantly affected by other mutual incompatibilities of the two parts of

the self incompatible solvent. Thus, we may limit us in this module to discuss the energy of

mixing only. Figure 8 shows schematically a thermodynamic cycle. The process we are

interested in (the addition of compound 3 to the preexisting mixture of compound 1 and 2) is

indicated by continuous arrows. The enthalpy of this process may be calculated from the other

two processes in this cycle; the preparation of a binary mixture of compound 1 and 2 and the

preparation of a ternary mixture from all three compounds in pure state.

123

1 2 3

12

233213311221123321 kkkVHmix

12*2

*123

*213

*1

33 1123312 kkk

VHmix

12*2

*1

233213311221

123312

kV

kkkV

H

o

mix

+

12*2

*1*

1221 kVHmix

123

1 2 3

12

233213311221123321 kkkVHmix

12*2

*123

*213

*1

33 1123312 kkk

VHmix

12*2

*1

233213311221

123312

kV

kkkV

H

o

mix

+

12*2

*1*

1221 kVHmix

Figure 8: Mixing thermodynamics of binary and ternary mixtures if all compounds are liquids.


21

The first of these two terms we already know from chapter 1.

12*

2*

1*1221 k

VHmix eqn. 40

Note that in this equation V and i are labeled with an asterisk. It must be taken into

account that the total volume of the binary mixture and its composition differs from the

ternary mixture. Thus we have to denote it with different symbols.

In the case of ternary mixtures, applying the same assumptions that were already

discussed in this chapter above lead to the following expression for the heat of mixing:

333222111233213311221332322

2211

21

123321 222 cccccccccV

Hmix eqn. 41

332322323313113122121121123321 222 cccccccccVHmix eqn. 42

Appropriate substitution: 12221211 2 kccc , 13331311 2 kccc , 23332322 2 kccc ,

yields the main simple eqn. 43 :

233213311221123321 kkkVHmix eqn. 43

The most interesting effect arises when one compound (e.g. molecules of type 3) is added into

a preexisting mixture of two compounds (e.g. molecules of type 1 and 2), since performing

the whole cycle has to yield in total a zero change in enthalpy, one obtains:

21*122

*2212

*211

*11333

233213311221332322

2211

21

123321 222

ccccc

ccccccV

Hmix eqn. 44

The reduction of eqn. 44 is given by

12*2

*13233213311221

123321 1 kkkkV

Hmix eqn. 45

The first three terms describe the energy of mixing of the direct preparation of the ternary

mixture. The last term is the energy of mixing that is needed to prepare the binary mixture of

compounds 1 and 2 of the original volume fractions *1 and *

2 , which is released upon the

conversion of the binary mixture into the ternary one. Note that this contribution to the

energy of mixing is negative. Given the fact that 1 and 2 are equal to *131 and

*231 respectively, the equation above can further be simplified to

12*2

*123

*213

*1

33 1123312 kkk

VHmix eqn. 46


22

The term in square brackets depends only on the constant factors k12, k13, k23 and the

composition of the preexisting mixture but is independent on the volume fraction of

compound 3. The denominator in the left term contains the term 33 1 which is identical

to the corresponding terms already shown in binary mixtures. The most important fact in eqn.

46 is that the term in square brackets can become negative, provided the parameters k13 and

k23 are small compared to the k12. Thus, even in systems comprising only symmetric

interactions, negative energies of mixing may be realised. However, a large value of the

parameter k12 is equivalent to the fact that a physical binary mixture of the two compounds 1

and 2 can not be achieved. The concept of the contribution presented here is to take

compounds 1 and 2 with a large interaction parameter k12 and to link them together by a

covalent bond to the molecule that is called, as described before: Self Incompatible Solvent.

This finally should give rise to a low or even negative energy of mixing with the third

compound.


23

However, if there are not only liquids mixed but the third substance is crystalline the

heat of fusion ( 3Hfus ) need to be taken into account as an additional term. This term

describes the entropy and energy participation needed for breaking the crystalline structure.

This term is shown in Figure 9 in a cyan coloured frame.

+

123

1 2 3

12

3233213311221123321 HkkkVH fusmix

12*2

*123

*213

*1

33 13123312 kkk

VHH fussol

12*2

*1*

1221 kVHmix

312*2

*1

233213311221

123312

HkV

kkkV

H

fuso

mix

+

123

1 2 3

12

3233213311221123321 HkkkVH fusmix

12*2

*123

*213

*1

33 13123312 kkk

VHH fussol

12*2

*1*

1221 kVHmix

312*2

*1

233213311221

123312

HkV

kkkV

H

fuso

mix

123

1 2 3

12

3233213311221123321 HkkkVH fusmix

12*2

*123

*213

*1

33 13123312 kkk

VHH fussol

12*2

*1*

1221 kVHmix

312*2

*1

233213311221

123312

HkV

kkkV

H

fuso

mix

Figure 9: Mixing thermodynamics of binary and ternary mixtures if one compound iscrystalline.

At this point the complete description of the thermodynamics of mixing for binary and ternary

mixtures is obtained. In principle, the parameters kij may be obtained from calorimetry and

binary mixing experiments. In Figure 9 the parameters in the red framed equation like:

volume fractions can be chosen by the experimentalists at will, the enthalpies of fusion are

known, can as well be obtained from calorimetric experiments or even better they are already

published for a lot of substances. The enthalpies of mixing can be determinate by solution

calorimetric experiments (see section 7.3). This gives rise to the experimentally defined ksis-3

values.


24

2.3 The strategies for experimental verificationIn the previous section it is developed a quantitative module that predicts a correlation

between heats of mixing and interaction parameters, that in principle all can be measured. The

only thing which is still needed is an intuitive graphical representation that allows to judge,

whether the experiments actually reflect the predicted trend: especially to judge whether the

negative term in the square brackets can be confirmed. For simplicity, the term on the left side

of the equation in the red frame in Figure 9 is called from now ksis-3. Furthermore, it is

assumed that the volume fractions of the preexisting mixture are not variable for example they

might be given by the requirement of the stoichiometry mixing ratios. As a consequence we

are left with four interactions parameters:

1. Interaction between compound 1 and 3 k13


3. Interaction between compound 3 and self incompatible solvents ksis-3


Four dimensional correlations of all these parameters would be difficult to show in one clear

and easy understandable graph. The interaction parameters k13 and k23 have the same character

thus, they are added term up and use as sum further on. If this is done, a 3-dimensional graph

of the interaction parameters can be prepared (Figure 10). On the abscissa to the right the sum

of first and second interaction parameters is plotted. On the ordinate the third interaction

parameter, and on the abscissa pointing forward the fourth interaction parameter are plotted.

Figure 10: Postulated 3-dimensional correlation with thermodynamically defined values.

23*213

*1 kk

12*2

*1 k

333

3

1312

sisfusmix k

VHH

Experimental dataTheoretical plane

23*213

*1 kk 23

*213

*1 kk

12*2

*1 k

333

3

1312

sisfusmix k

VHH

Experimental dataTheoretical plane


25

In Figure 10 the tilted gray plane symbolized the expectation. In this illustration a positive

value means an endothermic process and a negative value an exothermic process. On the

abscissa to the right the sum of the positive terms in the square brackets of eqn. 47 (k13 + k23)

is shown. On the ordinate is shown the interaction parameter ksis3. This is exactly a parameter

which needs to become as low as possible. It is postulated that favorable interactions between

compound 1/compound 3 (active ingredient) and between compound 2/compound 3 (active

ingredient) lead to favorable interaction between Self Incompatible Solvent/compounds 3

(active ingredient). Since favorable interactions manifest themselves in low values of kij the

gray plane leans to the left. The most interesting dependency is shown on the abscissa

pointing forward which represent the negative term in the square brackets in eqn. 47. It is

postulated that the unfavorable interactions between compound 1/compound 2 lead to

favorable interaction between Self Incompatible Solvent/compounds 3 (active ingredient).

Because of this, the expected gray plane is leaning forward. The theoretical predictions for the

principle of Self Incompatible Solvent described in this chapter, shows that the grey plane in

the Figure 10 is mathematically completely described by the following equation:

12*2

*123

*213

*1

33

3123312

1kkk

VHH fussol eqn. 47

They can be drawn exactly without any experiment. To compare these predictions with

experiments a series of compounds may be taken subsequently as compound 1 and compound

2, the parameters k12, k13, k23 and ksis3 may be determined experimentally, the corresponding

data for each of the combination of the compounds may be drown at appropriate positions into

the diagram and compared to the theory without any fitting procedure being necessary.

ksis3 ,ordinate

abscissapointing forward

abscissa tothe right

ksis3 ,ordinate



ksis3 ,ordinate



Chapter 3Model systems

26

Chapter 3:Model Systems

Abstract

A system made out of various compounds that are used as compounds 1, 2 and 3 in a series of

calorimetric experiments is chosen in such a way that it allows as close as possible

comparison to the theory and as well allows extrapolation to the Self Incompatible Solvent in

which parts resembling compound 1 and compound 2 are covalently connected. The

theoretical prediction for the first model is verified for simple mixtures and compared to the

theory introduced in the Self Incompatible Solvent. The theoretical model is intended to help

understanding solvents that comprise two incompatible parts that are linked together

covalently. Both model systems show reasonable agreement with the theory. In addition, all

results will be discussed and an outlook that describes possibilities of further work on this

area will be presented.


27

3.1 Model substances

3.1.1 Active ingredients: Compound 3

In various applications, especially in pharmaceutics and agrochemical applications, one needs

to dissolve or formulate a chemical that is responsible for the desired effect, but is effective

only if applied as part of a suitable formulation. Such substances are often called “active

ingredients”. In order to suit this choice of wording, we will from now on call compound 3 as

well “active ingredient”. However, for the sake of proving the mathematical model detailed in

the chapter before, we do not need to use a pharmacologically active (and thus toxic and

expensive) chemical. We chose, instead to use three compounds which all have a high

tendency to crystallize due to extended Pi-systems (as “real” active ingredients often have)

and in addition can be considered “neutral”, “acid” and “basic”. Thus, the model substances

are of our choice (Table 2).

Table 2: Active ingredients used in this work.

Compound 3 - active ingredient

OH O

Anthracene 9-Anthracene carboxylic acid

N

Acridin

Compound 3 - active ingredient

OH O

Anthracene 9-Anthracene carboxylic acid

N

Acridin

All active ingredients are bought from Merck Company with the purity of 96%. In some of

the experiments the applied amount of active ingredient is only partially dissolved. We took

this into account during the determination of the final concentration in solution via

spectroscopy. This commercially available substances contain: anthraqinone, anthrone,

carbosole, fluorine, 9,10 – dihydroanthracene, tetracene or bioanthyl carbozole[21] as main

impurities. However, impurities that has higher solubility than the active ingredients might

significantly interfere with this approach and this need to be eliminated. Because of this all

active ingredients needed to be purified; anthracene is recrystalized from cyclohexane and

chromatographed on aluminium oxide 90 active neutral; 9-anthracene carboxylic acid is

recrystlized from ethanol and acridine is purified in a two steps procedure. In the first step,

acridine is recrystallized from n-heptane. In the second step, acridine is chromatographed on

activated charcoal, aluminium oxide 90 active neutral and then again recrystallized from a

ethanol/water mixture. The purification processes for all active ingredients are described in

details in the experimental part (section 7.1).


28

3.1.2 Model substances: Compounds 1 and 2

In this work, the model substances for the compound 1 and 2 are delivered from chemical

providers or if this was not possible, they were synthesized by me. As it is mentioned in

chapter 2, the principle of self incompatibility should show the most pronounced effect for

pairs of substances that “don’t like” each other, that means that they show unfavorable

interaction. To show the increase of unfavorable interaction between various pairs of

substances, it is decided to start with “similar” pairs, where unfavorable interaction are not so

significant, for example compound 1 and 2 contain alkyl chains with the same length (see

schema 1 and 2 in Table 3). The couples of substances described in the previous section are

bond together via covalent bounds. The compound 1 and 2 are linked together via ester group.

Molecules obtained after esterification consist of these two connected incompatible sides and

as it is mentioned in chapter 2, they are called Self Incompatible Solvent. In Table 3 are shown

Self Incompatible Solvents which are bought or synthesized. Detailed synthesis procedures are

given in the experimental part (chapter 7).


29

Table 3: Model substances for compound 1, compound 2 and Self Incompatible Solvent.Nr.

1O

O

O

O

O

O

2

O

O

O

O

O

O

3O

O

O

O

O

O

4O

O

O

O

O

O

5O

O

F

F

F O

O

O

O

F

F

F6

O

O

F

F

FO

O

O

O

F

F

F7

O

O

FF

F FF

FO

O

O

O

FF

F FF

F

8O

O

O

O

OO

OO

O

OO

O

9O

O

O

O

OO

O

O

O

OO

O

10O

O

O

O

OO

O

O

OO

11O

O

O

O O

O

O O

12O

O

O

O

N

O

O

O

N

O

13O

O

O

O

O

O

O

O

O

O

14O

O

O

O

O

O

Compound 1Self Incompatible Solvent

linkerPart 1 Part 2Compound 2


30

3.1.3 Strategies of getting results

All substances, which are shown in Table 2 and Table 3, are used for the Precision Solution

Calorimetry experiments, to obtain the enthalpy of mixing. The detailed procedure of one

Precision Solution Calorimetry experiment is described in the experimental part. All solution

calorimetry experiments in this work are made at approximately 25°C and with a constant

amount of the active ingredients, which are enclosed in the glass break ampoule. Every

measurement is repeated at least two times. Initially the ampoules were closed with

commercially available glue “UHU plus finish solid 300” but during the progress of the thesis

it was recognized that a reaction between the glue and the measured substances occurred

because the baseline was not constant, see Figure 11.

Figure 11: Temperature profile of a calorimetric experiment if a reaction between the glueand closed substances occurs (baseline is not constant).

To obtain “flat” baseline several types of commercially available glues like: “UHU glass

reparation glue”, “Pistol Hot Pattex (Henkel)”, “UHU plus fast solid” were tested. Due to the

fact that all of them reacted with closed substances in ampoules, a droplet of molten glass (see

Figure 12) is used to close the ampoules. Integrity was checked by filling an ampoule with

anhydrous cobalt nitrate (blue) and immersing it into water (see Figure 12 a-c). Even after few

months no color transition was observed.

Methylpivalat in Tertbutylacetate

-4000000

-3000000

-2000000

-1000000

0

1000000

2000000

3000000

0 200 400 600 800 1000 1200 1400 1600 1800 2000

Time (s)

Tem

pera

ture

(K) Baseline

Baseline

Baseline

BaselineBreak

Calibration Calibration


31

a.

d.c.

b.

5cm 2.5cm

0.75cm 3cm

a.

d.c.

b.a.

d.c.

b.a.

d.c.

b.

5cm 2.5cm2.5cm

0.75cm0.75cm 3cm3cm

a) b)

c) d)

a.

d.c.

b.

5cm 2.5cm

0.75cm 3cm

a.

d.c.

b.a.

d.c.

b.a.

d.c.

b.

5cm 2.5cm2.5cm

0.75cm0.75cm 3cm3cm

a) b)

c) d)

Figure 12: Ampoules closed with molten glass and filled with: a-c) anhydrous cobalt nitrate (blue) and immersed in water d) anthracene.

When the ampoules are closed with molten glass the baseline was constant, see Figure 13.

Figure 13: Temperature profile of a calorimetric experiment if the ampoule is closed witha droplet of molten glass (baseline is constant).

The tightness of ampoules using this approach was proven and the first series of

measurements was conducted. However, reliability of the sealing still needed improvement

because often ampoules were broken after hot glass deposition onto the ampoule. Finally,

another commercially available two-component epoxy glue: “UHU plus directly solid” was

found. This glue had to be solidified for 1 day on the ampoules, after this period of time it was

successfully used in the solution calorimetry experiments (the baseline was constant). The

Methylpivalate in Tertbutylacetate

-4000000

-3000000

-2000000

-1000000

0

1000000

0 200 400 600 800 1000 1200 1400 1600 1800 2000 2200

Time (s)

Tem

pera

ture

(K)

BaselineBaseline

Baseline Baseline

CalibrationCalibration

Break


32

second series of measurements were done with ampoules closed with glue “UHU plus directly

solid”.

After some calorimetric experiments it is observed that the active ingredient is only

partially dissolved. This partial solubility gives rise to a heat of solution that depends on the

saturation solubility and may be significantly lower than “usual” raw data. To calculate

specific heats of dissolution, the actually obtained final concentration could not be determined

by dividing the amount by the total volume of the vessel, but needed to be evaluated

independently. The amount of active ingredients that are dissolved was measured via

fluorescence spectroscopy and then the values of mixH [kJ/mol] were calculated (see

appendix 1). The values of mixH that are calculated by determining the dissolved amount

from fluorescence measurements are marked in Table 9, Table 10 and Table 11 (chapter 7)

with an asterisk.


33

3.2 Physical 50/50 mixtureIn a first series of experiments a system as close as possible to the theory was used. The

compounds 1 and 2 had interaction parameters low enough that they still formed a

homogenous physical mixture. For a schematic illustration of the measurements performed to

obtain one data point, see Figure 14.

physical 50/50 mixture

neutra

llyneutrally

unfavorableinteraction

50 50

compound 1 compound 2

compound 3

compound 1

compound 2


neutra

llyneutrally


50 50


compound 3

compound 1

compound 2

Figure 14: First model system: physical 50/50 mixtures.

In the first series of experiments the enthalpy of mixing of compound 1 and 2 is measured and

interaction parameter k12 is calculated. In the next step compound 3 is added to the pure

compound 1 and from the enthalpy of mixing the interaction parameter k13 is calculated. The

same procedure is done with compound 2 to obtain interaction parameters k23. Finally a

homogeneous solution of compound 1 and compound 2 is prepared. These compounds are

mixed in the ratio 50:50 by volume. Then, the enthalpy of mixing, obtained after addition of

compound 3 into the mixture of compounds 1 and 2 is measured. From this experiment the

interaction parameter k50/50 mixture-3 is calculated. The calculation procedure is described in

details in chapter 2 and in the appendix 1) For these measurements are taken substances which

are shown in previous section (Table 2 and Table 3). One example for the model physical

50/50 mixture with chemical substances is shown in Figure 15.


34


neutra

llyneutrally


5050

O

O

O

O

N

O

O

O

O

O

N

O

HO O

k12

k13 k23

k50/50 mixture -3


neutra

llyneutrally


5050

O

O

O

O

N

O

O

O

O

O

N

O

HO O

k12

k13 k23

k50/50 mixture -3

Figure 15: First model system: physical 50/50 mixtures with chemical substances.

After the calculation of the interaction parameters for all mixtures, a three dimensional

diagram of 350/5033 1

123312mixtures

fusmix kV

HH (ordinate pointing upwards) vs.

23*213

*1 kk (abscissa pointing to the right) and 12

*2

*1 k (abscissa pointing forward) is

prepared.

3.2.1 ResultsThe obtained results are shown in a three dimensional diagram as discussed in the

previous chapter (Figure 10). Into this 3-dimensional plot of

350/5033 1

123312mixtures

fusmix kV

HH (ordinate pointing upwards) vs. 23

*213

*1 kk (abscissa

pointing to the right) and 12*2

*1 k (abscissa pointing forward) the data points corresponding


35

to each of the triplets are entered and the positions of all data points are compared to the

theoretical expectation of 12*2

*123

*213

*1

33 1123312 kkk

VHmix , see Figure 16.

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420

k 50/

50 m

ixtu

res-

3[1

03 JL-1

]

k 12

* 1* 2

[103

JL-1

]k13 *

1 + k23 *2 [103JL-1]

Figure 16: 3-dimensional correlation of the results from model: physical 50/50 mixture.Yellow points: measurements with anthraceneCyan points: measurements with acridineViolet points: measurements with 9-anthracene carboxylic acidGreen plane: fitted to the results from physical 50/50 mixtureRed plane: theory planeNumbers in frames indicate the model substances according to Table 3.

The theoretical prediction according to eqn. 47, without any fitting parameters is

indicated by the red frame. Yellow, cyan and violet points indicate measurements for active

ingredients; anthracene, acridine, 9-anthracene carboxylic acid respectively. Numbers in

frames in this diagram corresponds to numbers of schemes; see Table 2 and Table 3. It can be

12

13

1089


36

seen that there is a good correlation between the experimental data and the theoretical

prediction (red frame). In this 3-dimensional graph exact comparison is hindered by the fact

that actually only a 2-dimensional projection can be seen. Thus, to get a better impression of

the deviation, in addition a tilted plane is fitted to the experimental data and is represented as

a green frame (in Table 4 are data for the fitted green plane). Both frames are close to each

other indicating good agreement between theory and experiment. The frames are tilted to the

left and the most important fact is that the both planes are also tilted forward, as it is expected

from the theory.

Table 4: Data for the frame fitted to the results obtained from 50/50 physical mixtures.

Z = a + bX + cY

Z=-144668 + 0.9879X – 1.08604Y

a = -144668 error = 4191

b = 0.9879 error = 0.06026

c = -1.08604 error = 1.1441

Standard Deviation = 17.994 R2 = 0.941


37

3.3 Self Incompatible Solvent As already discussed in order to maximize the desired effect, systems in which

compound 1 and compound 2 have to be forced together by a covalent bond are most

promising. In the section 3.2 the model physical 50/50 mixture proved the thermodynamical

prediction. It would be interesting to check if there is also such a correlation when compound

1 and 2 are forced to exist together in one molecule, where compound 1 and 2 are connected

via a covalent bound. Still it is possible to generate k12, k13, k23 using compounds 1 and 2

which are as close as possible resemble the two parts of the self incompatible solvents. In the

second series exactly this was done. For that reason a second model system called Self

Incompatible Solvent is created. The schematic illustration of this model is shown

in Figure 17.


neutra

llyneutrally


linkerpart 1 part 2


compound 3


neutra

llyneutrally


linkerpart 1 part 2linkerpart 1 part 2


compound 3

Figure 17: Second model system: Self Incompatible Solvent.

In order to conduct a qualitative comparison between theory and experiment in such systems

it was the problem that a covalent bond disturbs the system. There is no way to sever or

tighten a covalent bond between compound 1 and compound 2 without influencing the

molecular properties. Thus in a purist approach, either separate compound 1 and 2 which are

needed to determine k12, k13, k23 or the covalently connected self incompatible solvent, made

exactly out of compound 1 and 2 are not available. There is only one difference compared to


38

the physical 50/50 mixtures model. In this model the enthalpy of mixing of compound 3 is

measured not in a physical mixture of compound one and two but in one molecule - Self

Incompatible Solvent - where the compound one and two are linked together via a covalent

bond. In this case, to obtain the Self Incompatible Solvent the alcohol component of the first

compound and the carbonic acid component of compound two are linked together by

esterification. The volume fraction of the two linked parts of the Self Incompatible Solvent,*1 and *

2 are estimated from molecular models. The energies of mixing of the compounds 3

with the Self Incompatible Solvents are determined. One example of such model system is

shown in Figure 18.

neutra

llyneutrally


O

O

O

O

N

O

HO O

k12

k13 k23

ksis-3


O

O

N

O

neutra

llyneutrally


O

O

O

O

N

O

HO O

k12

k13 k23

ksis-3


O

O

N

O

Figure 18: Example for the model system: Self Incompatible Solvent.

The same as in the case of the physical 50/50 mixture model the enthalpies of mixing for Self

Incompatible Solvent are measured with three sorts of compound 3: anthracene, acridin and

9-anthracene carboxylic acid. Afterwards the interaction parameters k12, k13, k23, ksis3 are

calculated (see appendix 1).

3.3.1 ResultsAs already mentioned in the previous section, the same examples of compounds 1 and

2 as for the physical 50/50 mixture model system are used for the Self Incompatible Solvent


39

model system with the same three active ingredients (compound 3). In this case compounds 1

and 2 are linked together and form one molecule. After determination of all the interaction

parameters the corresponding diagram (see Figure 19) is created. Then the positions of all

experimental data are compared to the theoretical expectation of

12*2

*123

*213

*1

33 1123312 kkk

VHmix .

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[1

03 JL-1

]

k 12

* 1* 2

[10

3 JL-1

]

k13*1 + k23

*2 [103JL-1]

Figure 19: 3-dimensional correlation of results from the model system: Self Incompatible Solvents.

Yellow points: measurements with anthraceneCyan points: measurements with acridineViolet points: measurements with 9-anthracene carboxylic acidBlue plane: fitted to the results from Self Incompatible SolventsRed plane: theory planeNumbers in frames indicate the model substances according to Table 3.

11139

10

12


40

In Figure 19 a planar fit of the experimental data is indicated by the blue plane (data for the

fitted blue frame are shown in Table 5). The theoretical prediction without any fitting

parameters is again indicated by the red frame. Both planes are tilted to the left. The most

important fact in this diagram is that both planes are tilted forward, as expected by the

theoretical predictions in chapter 2. That means that the self incompatibility of the solvents

improves the dissolving process. There is a systematic deviation between theory and the

experimental plane.

Table 5: Data for the frame fitted to results obtained from Self Incompatible Solvents.

Z = a + bX + cY

Z= -51438 + 0.67987X – 3.1848Y

a = -51438 error = 7195

b = 0.67987 error = 0.10427

c = -3.1848 error = 1.982


Given the simplicity of the model the discrepancy is moderate. It is most likely due to the fact

that in the theory compound 1 and compound 2 are identical to the two parts liked together in

the Self Incompatible Solvent. However, in this case, the model compounds in addition

comprise the acetoxy and the methoxy - group respectively.

3.4 Conclusion and outlook

The two models systems: physical 50/50 mixture and Self Incompatible Solvent are

described and examined. In order to compare both systems, compounds are chosen which are

completely miscible, even as separate molecules. In both 3-dimensional diagrams the data for

the 50/50 mixture of two compounds and for the Self Incompatible Solvent are represented by

a plane that is tilted forward. The thermodynamic principle described in chapter 3 is working.

An unfavorable interaction between compound 1 and compound 2 is favorable for the

dissolution of a third compound, which means that the self incompatibility of the solvents

improves the dissolution process. The agreement between theory and experimental data is

much better for the system physical 50/50 mixture than for the covalently linked Self


41

Incompatible Solvents. This fact is to be expected, because the introduction of the covalent

link is an additional modification of the system and this change is not taken into account by

the theory. These results promise more drastic effects by compounds that are so incompatible,

that they would not form physical mixtures by themselves but must be forced together by

covalent linking. These results suggest that also by such linked system some energy can be

won for dissolving third compound.

To show more drastic effects of the Self Incompatible Solvent, a new solvent molecule

can be prepared which would consist of: for example long alkyls chain on the left side and

ionic liquid molecule on the right side. This concept can be also transfer from low molecular

mass substances to polymers. In this case, polymer chains should consist of for example one

hydrophilic and one hydrophobic part.

Chapter 4Mixing thermodynamics of binary mixtures

42

Chapter 4:Mixing thermodynamics of binarymixtures

Abstract

By the usage of the standard equipment for the Precision Solution Calorimeter it is possible to

measure the energy of mixing of two compounds, but only if one compound is in excess to the

second one. This chapter describes a method to obtain the enthalpy of mixing of two

compounds at any desired mixing ratio, using the same equipment. Additionally, experimental

results are compared to the values of the enthalpies obtained from the theoretical predictions

of Van Krevelen and Hoy.


43

4.1 Compound 1 in Compound 2

In Figure 20 are shown theoretical enthalpies according to mean field theoryV

Hmix versus

volume fraction of one compound (blue, dash dotted and dotted lines). Heats of dissolution

can be measured with solution calorimetry. In this technique one compound is enclosed

within a sealed glass ampoule immersed into the other compound. After thermal equilibration,

the glass ampoule is broken and both compounds are mixed. The change in temperature is

recorded and converted into the heat of dissolution using the heat capacity of the filled

container. However, glass ampoules have a volume of 1 ml, which is significantly smaller

than the volume of the reaction vessel (25 ml). Thus, it is only possible to measure binary

mixtures, if one compound is in excess to the other one. The predicted dependency is

experimentally verified over the entire range of the composition. Thus, within the hatched

region in Figure 20, no data can be generated using this conventional approach.

Figure 20: Theoretical enthalpy of mixing plotted versus volume fraction of one of the mixedcompounds.

To obtain experimental results also from the hatched area, one can design a thermodynamic

cycle and devise the desired date by applying the Hess’s law.


44

4.2 Dilution experiment and Hess lawThe Hess’s law states that: "The standard enthalpy of an overall reaction is the sum of

the standard enthalpies of the individual reactions into which a reaction may be divided[1]". In

this work the Hess’s law is used for the dissolution experiments and finally for the calculation

of the enthalpies of mixing of two compounds in a ratio smaller than 1:25. Thus, if a chain of

events, a so called thermodynamic cycle, finally restores the initial conditions of a system, no

heat of reaction occurs. If a thermodynamic cycle is performed in all but one of the events, the

sum of the heat of reactions of all these events is equal to the corresponding value of the

missing event. Thus, an appropriate thermodynamic cycle was designed, in which all events

were measurable except to one that is interested. This cycle is shown in Figure 21. Actually,

Figure 21 depicts two independent and twined cycles (indicated by cyan dotted and violet

dashed lines). Only one of them is needed, for symmetry reasons both are shown. The desired,

but experimentally inaccessible experiment (mixing 1 and 2) is indicated in red solid line.

Since this is the only event that is inaccessible the remainder of the cycle can be used to

obtain the desired value.

1 2

1+2

Solvent

)22( SSmix H

2 + Solvent

)1221( SSmix H1+2 + Solvent

)1212( SSH

1 + Solvent

)11( SSmixH

)1212( SSmix H

1 2

1+2

Solvent

)22( SSmix H )22( SSmix H

2 + Solvent

)1221( SSmix H )1221( SSmix H1+2 + Solvent

)1212( SSH

1 + Solvent

)11( SSmixH )11( SSmixH

)1212( SSmix H )1212( SSmix H

Figure 21: Dissolution experiment.

First the enthalpy of mixing of compound 1 in the solvent (here: acetone) and then the

enthalpy of mixing of compound 2 in this solution of compound 1 in solvent is measured.

Next, the same amounts of pure compound 1 and 2 as were used for the previous


45

measurements are mixed together. Afterwards the enthalpy of mixing of this mixture with the

same solvent is measured. Below, there is given one example for the enthalpy of mixing of

two compounds mixed in the 1:1 ratio:

1.) 0.4 ml of compound 1 (closed in a small glass ampoule) is added during the break

experiment into 25 ml of the solvent ( SSmix H 11 )

2.) 0.4 ml of compound 2 (closed in a small glass ampoule) is added during the break

experiment into the mixture of 0.4 ml of compound 1 in 25 ml of the solvent

( SSmix H 1212 )

3.) 0.4 ml of compound 1 and 0.4 ml of compound 2 are mixed together and closed in the

small glass ampoule, then this mixture is added into the solvent ( SSmix H 1212 )

during the break experiment.

After these three measurements, the enthalpy of mixing of compound 1 and 2 in the 1:1 ratio

is calculated (see eqn.48).

SSmixSSmixSSmixmix HHHH 12121212111221 eqn. 48

A similar procedure can be used for the calculation of the enthalpy of mixing of compound 1

in 2 if it is started with compound 2. The schematic illustration of this procedure is shown on

the right side in Figure 21. Finally, the enthalpy of mixing is given by the equation

SSmixSSmixSSmixmix HHHH 12121221221221 eqn. 49

Using this method the enthalpies of mixing of compound 1 in 2 and vice versa are measured

and calculated for the following volume fractions: 0.2, 0.4, 0.5, 0.6 and 0.8.

From all examples of compounds 1 and 2 presented in the Table 3 (chapter 3), six schemes

are chosen to present the behavior of this compound in binary mixtures, see Table 6.


46

Table 6: Chosen examples for binary mixtures.Nr.

3O

O

O

O

10O

O

O

O

OO

11O

O

O

O O

12O

O

O

O

N

O

13O

O

O

O

O

O

14O

O

O

O

These schemes are chosen because compounds 1 and 2 differ from each other in their

chemical structures, but they are still miscible in the 1:1 ratio.

4.3 Results for binary mixturesThe method described in the previous section is used for the calculation of enthalpies of

mixing of two compounds that are mixed for the following volume fractions: 0.2, 0.4, 0.5, 0.6

and 0.8. In a first series of experiments acetone is used as solvent. The values that are finally

obtained are small differences between large numbers. This finally causes comparatively large

experimental errors. To minimize these errors, all microcalorimetry measurements are

repeated three times and then the average value is calculated. In the following two diagrams

(Figure 22 and 23) all data points with the exception of the two most outwards ones were

obtained using Hess law. The most outwards ones were obtained in a conventional break

experiment with the minority of compound being inside the break ampoule. The Figure 22

nicely shows the expected cone shape of the data. However, the outmost data points are all

very close to the abscissa and relative errors in these points might be underestimated. In order

to reveal that these points as well agree with the expected line in Figure 23 1221

kV

Hmix is

plotted on the ordinate as a function of 1 .

Compound 1 Compound 2


47

Figure 22: Enthalpy of mixing via volume fraction;blue points: schema 12, dark green points: schema 10, red points: schema 11cyan points: schema 13, green points: schema 14, brown points: schema 3

Figure 23: Enthalpy of mixing divided by volume fraction as function of the volume fraction;blue points: schema 12, dark green points: schema 10, red points: schema 11cyan points: schema 13, green points: schema 14, brown points: schema 3


48

In Figure 23 it is expected, that all data points corresponding to one scheme to be aligned on a

horizontal line. This indeed is the case, thus there is agreement between both methods. As it

is expected, for a given scheme, the values of 12k for various volume fractions of mixed

substance are identical within the experimental errors. This is indicated in Figure 23 as dotted

lines. The heights of the dotted lines in Figure 23 correspond to the difference in the chemical

structures of the compounds used for the experiments. The values 12k for chosen pair of

substances are in good agreement with the structural difference within that pair. With increase

of the differences in chemical structure between both compounds, the increase of the value

12k is observed. Both compounds from schemes 3 and 14, (see Figure 24) are unpolar; the

interactions between these two compounds are favorable and this cause small values of 12k .

The biggest value of 12k , which is equivalent to the strongest unfavorable interaction, is

obtained from the schema 12, where one compound is unpolar - ester with long alkyl chain

and second compound is polar - contains pyrolidon in the chemical structure.

O

O

O

O

14

3

13

10

12

Incr

easi

ng k

12

O

O

O

O

O

O

O

O

O

O

O

O

O

O

N

O

11

O

O

O

O O

O

O

OO

O

O

O

O

O

O

14

3

13

10

12

Incr

easi

ng k

12

O

O

O

O

O

O

O

O

O

O

O

O

O

O

N

O

11

O

O

O

O O11

O

O

O

O O

O

O

OO

O

O

Figure 24: Schemes with increasing value 12k from bottom to top.


49

4.4 Results with various solventsIn chapter 4.2 is described the method of the determination of the enthalpy of mixing of two

compounds by using acetone as solvent. Acetone is a polar aprotic solvent with a dipole

moment of 9.54·10-30 [Cm], according to Hess’s law the obtained values are expected to be

independent on the chosen solvent. However some might be more reliable than others. It is

interesting to check, if the chosen method is independent of the solvent character and its

polarity. Therefore the enthalpies of mixing of all pairs of substances that are shown in

section 4.2 and 4.3 are measured in five various solvents with different characters and

polarities. The following solvents are chosen for the measurements:

Methanol - polar protic solvents with dipole moment of 5.67·10-30 [Cm]

Ethanol – polar protic solvents with dipole moment of 5.77·10-30 [Cm]

Acetone – polar aprotic solvent with dipole moment of 9.54·10-30 [Cm]

Tetrahydrofuran (THF) – polar aprotic solvent with dipole moment of 5.84·10-30 [Cm]

n-Hexane – non polar solvent with dipole moment of 0.00 [Cm].

All values are taken from the textbook [22].

Similar as in case where acetone was used as solvent, each of the measurements with different

solvents are repeated three times to decrease the error. In Figure 25 are shown the enthalpies

of mixing of six pairs of substances, which were measured in five different solvents. Each

solvent is represented via two columns. Generally, the first column corresponds to the

thermodynamic cycle depicted by a cyan dotted line in Figure 21; the second column

corresponds to the violet dashed line shown in the same figure.


50

Methanol Ethanol Aceton nHexan THF0

2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12

12

O

OO

O

N

O


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12

12

O

OO

O

N

O


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 12

12

O

OO

O

N

O

12

O

OO

O

N

O

10

O

O

O

O

OO


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10

10

O

O

O

O

OO

10

O

O

O

O

OO


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10


2000

4000

6000

8000

Hm

ix[k

J/m

3 ] Sc

hem

a 10


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13

13O

O

O

O

O

O


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13

13O

O

O

O

O

O


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13


1000

2000

3000

4000

5000H

mix

[kJ/

m3 ]

Sch

ema

13

13O

O

O

O

O

O13O

O

O

O

O

O


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11

11

O

O

O

O O


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11

11

O

O

O

O O


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11


2000

4000

6000

8000

Hm

ix[k

J/m

3 ]Sc

hem

a 11

11

O

O

O

O O11

O

O

O

O O


1000

2000

3000

4000

5000

6000

Hm

ix[k

J/m

3 ]Sc

hem

a 14


1000

2000

3000

4000

5000

6000

Hm

ix[k

J/m

3 ]Sc

hem

a 14

14O

O

O

O


1000

2000

3000

4000

5000

6000

Hm

ix[k

J/m

3 ]Sc

hem

a 14


1000

2000

3000

4000

5000

6000

Hm

ix[k

J/m

3 ]Sc

hem

a 14

14O

O

O

O14O

O

O

O


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3

O

O

O

O3


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3


500

1000

1500

Hm

ix[k

J/m

3 ] Sc

hem

a 3

O

O

O

O3O

O

O

O3

Figure 25: Enthalpy of mixing for six schemas with five various solvents.

As it is expected, the character and the polarity of the solvent have no significant influence on

the enthalpy of mixing of the chosen compounds. On the other hand the statistical errors of

those data obtained with polar solvents are often significantly higher than the corresponding

values for non polar solvents. This most probably is due to the fact, that the heats of


51

dissolution in the former case are significantly larger. Thus the desired value thus is a small

difference between large numbers and thus less reliable.

4.5 Comparison of the experimental data with the existingtheories

Calorimetry allows the experimental determination of the parameters kij that describe

the pairwise interactions between two compounds. As already discussed in chapter 2, there are

schemes that allow estimating these pairwise parameters from other experimentally accessible

parameters that describe the properties of individual compounds, e.g. the Hildebrand or

Hansen parameters. An even further reduction in complexity is offered by the so called group

contribution methods e.g. van Krevelen or Hoy’s methods, that allow estimating the

parameters that characterize the individual compounds based on the chemically distinctive

groups that these compounds comprise.

In section 1.5 is shown the comparison between van Krevelen and Hoy theories, where

the cohesion energy density of the compound i due to "dispersion interactions" di , "polar

interactions" pi and "hydrogen bonds" hi may be estimated from group contributions. Both

calculation schemes are developed independently of each other and make similar predictions.

In general, for the calculation of di , the following procedure is used: di = a (number of

CH3-groups) + b (number of –OH – groups) + … . The comparison of van Krevelen and

Hoy methods gave results which show a rough correspondence, for details see section 1.5,

Figure 6. It is interesting to compare these theories with the results obtained from the solution

calorimetry. For the comparison, model substances for compound 1 and compound 2 (see

Table 3, section 3.1.2) were used. Afterwards the enthalpies of mixing are measured via

solution calorimetry.


52

The experimental data are then compared in the following equation:

eqn. 50

van Krevelen or Hoy experimental

The left side of the equation is computed by van Krevelen or Hoy (see chapter 1) methods and

the right side of the equation is obtained from the solution calorimetry measurements. In

Figure 26 the right side of equation is plotted versus the left side.

-50 0 50 100 150 200 250 300

0

2000

4000

6000

8000

10000

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y tw

o m

etho

ds

mixH / (kJ/m3)experimental results from solution calorimetry

-50 0 50 100 150 200 250 300 350-50

0

50

100

150

200

250

300

350

calculated by Hoy methodcalculated by Hoftyzer - van Krevelen method

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y tw

o m

etho

ds


-50 0 50 100 150 200 250 300

0

2000

4000

6000

8000

10000

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y tw

o m

etho

ds


-50 0 50 100 150 200 250 300 350-50

0

50

100

150

200

250

300

350

calculated by Hoy methodcalculated by Hoftyzer - van Krevelen method

D 1-

D 2)2 +(

P 1-P 2)2 +(

H 1-

H 2)2

(kJ/

m3 )

calc

ulat

ed b

y tw

o m

etho

ds


Figure 26: Comparison of the experimental data with Hoy’s and van Krevelen’s theories.

In Figure 26 can be seen that the experimental results from solution calorimetry show coarse

correlation with Hoy’s and van Krevelen’s theory. The biggest deviation can be seen by

compounds that contain fluor in their chemical structures. In the graph in the blue frame are

shown points from the area marked with blue dotted line. In this area the best correlation is

obtained. Although, the points in this area are in the same scale bars, this correlation is not

good enough. Due to the fact that the correlation between Hoy’s and van Krevelen’s theories

is not satisfying (see chapter 1), the results shown in Figure 26 are not surprising but

acceptable.

21

221

221

221

HmixHHPPDD

Chapter 5Saturation Solubility

53

Chapter 5:Saturation Solubility

Abstract

In the previous chapters the interaction parameter between active ingredients and Self

Incompatible Solvents was obtain from solution calorimetry. However, the same parameter

can also be achieved from the saturation solubility. In this chapter the theoretical description

of the extraction of this parameter will be described. Additionally, solubility experiment will

be explained and finally the comparison of the results obtained from the calorimetry and

saturation solubility will be shown.


54

5.1 Saturation solubilityThe goal of the investigation is to prove that the self incompatibility of the solvent

improves the dissolution of crystalline substances. In the chapter 2 the theoretical predictions

for Self Incompatible Solvents were verified with experimental data. To obtain these results, a

Precision Solution Calorimeter was used and the enthalpies of mixing were measured.

However, one may analyse the saturation solubility instead of the enthalpy of mixing, to

obtain the same parameters as before. This opens a second, independent approach. By theory

the both approaches shall give identical results.

Saturation solubility is achieved when the maximal amount of a substance is dissolved and an

equilibrium between undissolved and dissolved solute occurs. Such a solution is called

saturated and the chemical potential of the undissolved pure solid substance )s(¤3 is equal to

the chemical potential of the dissolved solute )l(3 . (In this work this substance is labelled as

compound 3, thus, its chemical potential is indicated with number 3.)

)](fln[RTlnRTlls 33¤33

¤3

eqn. 51

In Figure 27 are summarized graphics of entropy, enthalpy and the Gibbs free energy

of mixing from chapter 1. On the right ordinate of the graphics are shown three points; red,

blue and violet respectively. These points indicate the state of the pure crystalline substance

(compound 3). The composition of the saturated solution in equilibrium to the crystalline pure

substance (compound 3) is obtained by drawing a tangent to the curved line in diagram c

which in addition runs trough the isolated point on the right ordinate. The tangent touches the

curved line at the composition of the saturated solution. As can be seen from graph c, this

consideration depends on the interaction parameter ki3.

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

a) b) c)

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

a) b) c)

saturated solution

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

a) b) c)

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

0 0.2 0.4 0.6 0.81

0

1

2

10

VHsol

0 0.2 0.4 0.6 0.81

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

100 0.2 0.4 0.6 0.8

1

0

1

2

)

10

VSsol

VGsol

Incr

easi

ng k

i3

Solubility gap

a) b) c)

saturated solution

Figure 27: General diagrams of: a) the entropy; b) the enthalpy of mixing and c) the Gibbs free energy of mixing when one of mixed compounds is crystalline.


55

The mathematical description of the Gibbs free energy of mixing is shown in eqn. 52

33

3333 lnln

VG

kRTV

RTVV

G fusiii

isoleqn. 52

After mathematical transformation of eqn. 52 it is possible to calculate the relation between

the interaction parameter ki3 and the saturation solubility, (see next sections).

5.2 Extracting ki3 from saturation solubility5.2.1 Ideal solution

The solutions from the physical chemistry point of view are divided in ideal and

regular. To obtain a solution, two substances can be mixed, for example; a pure liquid

substance “i” (solvent), in which only “i-i” interactions occur with the pure crystalline

compound 3 (solute), in which only “3-3” interactions occur. As a result, a simple mixture is

obtained in which three

kinds of interactions arise;

(i) the interactions

between the pure

substance “i” (“i-i”), (ii)

the interactions between

the pure substance 3 (“3-

3”), and (iii) the

interactions between

substances “i” and

compound 3 (“i-3”). In

ideal solutions the average

energy of interactions in

the mixture “i-3”, “i-i”,

“3-3” is identical to the

average energy of interactions in pure substances: “i-i” and “3-3”. That means that the

interaction parameter ki3 is zero. If the molecules of the mixed substances have the same

molar volume, V the eqn. 52 for the Gibbs free energy of mixing is given by:

3

3

33 lnln

VG

RTV

RTVV

G fusi

isoleqn. 53

Figure 28: Schematic illustration of mixing two substances.

+i i

i 3i 3

3 3

i i

i 3

i 3

3 3

i i

3 3

i i

i i

i i

i i

ii

3 3

3 3

3 3

3 3

3 3

+i ii i

i 3i 3i 3i 3

3 33 3

i ii i

i 3i 3

i 3i 3

3 33 3

i ii i

3 33 3

i ii i

i ii i

i ii i

i ii i

ii

ii

3 33 3

3 33 3

3 33 3

3 33 3

3 33 3


56

The chemical potential is the first derivative of the Gibbs free energy with respect to the

molar amount of the considered compound n3. Therefore the change in the chemical potential

upon dissolution is given by (for a detailed derivation of the mathematical results given here

and below, see the appendix 2):

333

3 ln GRTdN

Gdfus

solsol

eqn. 54

At saturation, the chemical potential of the dissolved solute is equal to the chemical potential

of the undissolved solute, that means:

03sol eqn. 55

Thus, the logarithm of the saturation solubility in case of ideal solutions is proportional to the

Gibbs free energy of fusion of the crystal.

0ln 3,3 GRT fussat RTGfus

sat3

,3 expeqn. 56

This prediction reflects the common observation that the more difficult it is to melt a

substance the more difficult it is to dissolve it. In ideal solutions the interaction parameter is

per definition zero. However, ideal solutions, are seldom, thus it is necessary to include a non

zero value of ki3 into the description.

5.2.2 Regular solution

In case of regular solutions the interactions solute/solvent are not identical to the interactions

solute/solute and solvent/solvent. If the interactions solute/solvent are more favourable than

the interactions solute/solute and solvent/solvent, then dissolving process will be exothermic.

The second possibility is that the interactions solute/solute and solvent/solvent are more

favourable in comparison to the interactions solute/solvent. In this case the dissolving process

will be endothermic. Due to those facts, it is possible to extract the interaction parameter ki3 in

regular solution from the saturation solubility. When the derivation which above yielded eqn.

56 is applied, then the change in chemical potential upon dissolution, when both molecules

have the same molar volume, extracting ki3 from saturation solubility yields to eqn. 57:

32

333

3 ln GVkRTdN

Gdfusii

solsol

eqn. 57

and 0:saturationat 3sol


57

32,3

3,3

1ln

isat

fussat kV

GRT eqn. 58

When both molecules have different molar volumes one obtains:

332

3333

3 1ln GVVRTVkRT

dNGd

fusi

iiisol

sol

eqn. 59

and 0:saturationat 3sol

32,33

33

,3,3

1

1)1(ln

isat

fusi

satsat

kV

GVVRT

eqn. 60

Solutes used in this work (compound 3; active ingredients) and solvents (Self Incompatible

Solvent) have different molar volumes, therefore the interactions parameter ki3 from the

saturation solubility is calculated as it is described by eqn. 60.

It is worth noting that the change of the chemical potential upon changing the

concentration is usually given in the form of:

fRTRTfRT

aRT

lnlnlnln eqn. 61

The so called activity a, is given by the concentration times a concentration dependent

activity coefficient )(fa . Usually there are two considered cases: either the standard

potential is chosen in such a way that )(f =1 for the pure compound or in such a way

that )(f approaches the value 1 for the limiting case of an infinite solution. In each of both

cases the mathematical form of eqn. 61 is fulfilled. However, the values of the standard

potential and the activity coefficient are not identical. Similarly, one has the freedom to

express the concentration in terms of the volume fraction, the molar fraction, the molar

concentration etc. . This again will not cause deviations from the mathematical expression

but will influence the values that have to be used for the standard potential and the activity

factors )(f , )(xf , )(cf . To distinguish between these choices, one often denotes the

standard chemical potential of a description in which f approaches 1 at infinite dilution with

, while in a description in which the activity factor is 1 for the pure substance one usually

denotes the standard potential with µ .


58

Thus, if a description that uses the activity coefficients is preferred, the chemical

potential of substance three might be expressed in the form of:

3332

33333 111ln GVVRTVkRT fus

ii

crystalpure

eqn. 62

332

333333

333 111ln1i

ifusi

icrystalpure V

VRTVkRTGVVRTVk

eqn. 63

3 33ln fRT

As indicated in eqn. 63, the first group of terms of the right hand side can be interpreted as a

concentration independent standard potential, after it, a term that is the ideal concentration

dependence of chemical potential, while the last group of terms is the equivalent of

)(ln fRT .

5.3 Solubility experimentIn the previous section it was shown that the interaction parameter can be derived from

the simple eqn. 60. The left hand side of this equation comprises only parameters which can

either be estimated with reasonable accuracy or can be measured directly. iV is easily

calculated from the density and the molar mass, 3V was estimated from the density of the

molten compound (assumed to be 1g/cm3) and its molar mass and Gfus was obtained from

literature of the enthalpy of melting and its melting temperature. Only the saturation solubility

needed to be measured for each pair of compound i and compound 3. Each of Self

Incompatible Solvents (compound i), which are given in Table 3 and each of active

ingredients: anthracene, acridine, 9-anthracene carboxylic acid (compound 3) are mixed with

each other, in such a way that supersaturated solutions are obtained. Then all these solutions

are shaked at a temperature of 25°C for 4 weeks. After this time, the insoluble part of the

active ingredients is sedimented by centrifugation (25°C, 12000 rpm, 4 hours) and the

solution is decanted from the vessel. Next, the saturated solutions thus obtained, are diluted to

concentrations that allow a quantitative analysis via fluorescence spectroscopy. The

concentrations of the saturated solutions are calculated from the intensity and the dilution

factor and are used for the further calculation of ki3. In Figure 29 is shown the model system

of Self Incompatible Solvent already familiar from the Figure 18. Deviating from the Figure

18, the vertical arrow is now labeled ki3 instead of ksis-3. In principle both parameters should


59

be identical in this work, only different names are used to indicate the non identical methods

that were used to obtain them.


neutra

llyneutrally


linkerpart 1 part 2


compound 3

k12

k13 k23

ki3


neutra

llyneutrally


linkerpart 1 part 2linkerpart 1 part 2


compound 3

k12

k13 k23

ki3

Figure 29: Model system: Self Incompatible Solvents with ki3.

5.4 Results

5.4.1 Precision solution calorimeter versus saturation solubility

At this point, the interaction parameter between the Self Incompatible Solvent and active

ingredients is calculated in two different ways. In the chapter 3, this parameter is described as

ksis-3 and is calculated from the enthalpy of mixing obtained by using the Precision Solution

Calorimeter. In the previous section is shown the calculation of the same parameter from the

saturation solubility. In case of the calculation via saturation solubility this parameter is called

ki3. In principle, following the theory, the both parameters should have the same value for

each of used schemes, ksis-3 = ki3. To compare these results, a 2-dimensional graph is

prepared, where ksis-3 (from calorimetry) is plotted versus ki3 (from saturation solubility), see

Figure 30.


60

-200 -150 -100 -50 0 50-200

-150

-100

-50

0

50

acridin anthracene 9-anthracene carboxilic acid

k sis3

from

sat

urat

ed s

olut

ion

[MJ/

m3 ]

ksis3

from calorimetry [MJ/m3]

Figure 30: Interaction parameter ksis-3 (from calorimetry) versus interaction parameter ki3(from saturation solubility).

Following the theory, all points shown in Figure 30 should be located on the red

dashed line. Although, the values of ksis-3 and ki3 are in the same, the correlation is not

satisfying. The deviations might be a result of experimental errors. Although one might

consider, the procedure used to determine equilibrium solubilities used here as quite simple

and straightforward, it suffers from the difficulty to separate the solution from the non-

dissolved excess of the active ingredient. The mixtures usually contain very fine particles and

to some extend are quite viscous, thus it is difficult to exclude the possibility of incomplete

separation by the centrifugation process and thus concentrations might be overestimated.

The 3-dimensional comparison of data obtained from calorimetry and saturation

solubility are shown in Figure 31. In this 3-dimentional illustrations the data on the abscissa

pointed forward and on the abscissa pointing to the right; *2

*112k , *

223*113 kk respectively,

are identical. As already mentioned before, these data are calculated using measurements of

the enthalpy of mixing via Precision Solution Calorimetry. The ordinate in Figure 31a is the

value of ksis3: the interaction parameter between Self Incompatible Solvent and active

ingredients (compound 3) obtained from the solution calorimetry. The ordinate in Figure 31b

shows the value of ki3: the interaction parameter between the Self Incompatible Solvents and

active ingredients (compound 3) calculated from saturation solubility. As already mentioned,


61

in principle ki3 should be equal to ksis3. The theoretical prediction according to eqn. 47 without

any fitting parameters is indicated by the red frame, as in chapter 3. Yellow, cyan and violet

points indicate measurements for active ingredients: anthracene, acridine,

9-anthracene carboxylic acid, respectively. Numbers in frames corresponds to numbers of the

schemes in chapter 3 (Table 2 and Table 3). To get a better impression of the deviation, in

addition a tilted plane (see Table 7) is fitted to the data calculated from the saturation

solubility and is represented by a violet frame (Figure 31b). It can be seen that the results

obtained by the saturation solubility show worse correlation to the theoretical predictions than

the results obtained from calorimetry. The deviation is much stronger as in case of the model

system: Self Incompatible Solvent (comparison Figure 31a and b).

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[1

03 JL-1

]

k 12

* 1* 2

[10

3 JL-1

]

k13*1 + k23

*2 [103JL-1]

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[103 JL

-1]

k 12

* 1* 2

[103

JL-1

]

k13 *1 + k23 *

2 [103JL-1]

10

12

131110

9

b.)

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[1

03 JL-1

]

k 12

* 1* 2

[10

3 JL-1

]

k13*1 + k23

*2 [103JL-1]

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[103 JL

-1]

k 12

* 1* 2

[103

JL-1

]

k13 *1 + k23 *

2 [103JL-1]

10

12

131110

9

b.)a) b)

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[1

03 JL-1

]

k 12

* 1* 2

[10

3 JL-1

]

k13*1 + k23

*2 [103JL-1]

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[103 JL

-1]

k 12

* 1* 2

[103

JL-1

]

k13 *1 + k23 *

2 [103JL-1]

10

12

131110

9

b.)

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[1

03 JL-1

]

k 12

* 1* 2

[10

3 JL-1

]

k13*1 + k23

*2 [103JL-1]

-150 -120 -90 -60 -30 0 30

-150

-100

-50

0

50

8

6420k s

is-3

[103 JL

-1]

k 12

* 1* 2

[103

JL-1

]

k13 *1 + k23 *

2 [103JL-1]

10

12

131110

9

b.)a) b)

Figure 31: 3-dimensional correlation of results from a)calorimetry, Self Incompatible Solvent,

b)saturation solubility, Self Incompatible SolventYellow points: measurements with anthraceneCyan points: measurements with acridineViolet points: measurements with 9-anthracene carboxylic acidBlue plane: fitted to the results from calorimetryViolet plane: fitted to the results from saturation solubilityRed plane: theoretical plane

Numbers indicate the number of the model substances according to Table 3.


62

Table 7: Data for the frame fitted to the results obtained from the saturation solubility.

Z = a + bX + cY

Z= 15396 + 0.27119X – 6.48647Y

a = 15396 error = 5865

b = 0.27119 error = 0.08433

c = -6.48647 error = 1.601


Nevertheless, even if the interaction parameters between the Self Incompatible

Solvents and active ingredients are calculated from the saturation solubility, the violet plane is

also tilted to the left and the most important fact is that this plane is also tilted forward, as it is

expected from the theory. That means that if the solubility is considered, the principle seems

to work, the self incompatibility of the solvents favour the solubility of the active ingredients.

5.5 ConclusionThe interaction parameter between active ingredients (compound 3) and the Self

Incompatible Solvents can be obtained in two ways: from calorimetry and from saturation

solubility. The comparison of the results from these two methods showed a non uniform

correlation. The differences in the obtained data can be caused by the higher complicity of the

saturation solubility experiments. Even very careful work during these complicated

experiments could not eliminate some errors. On the other hand, the second reason for this not

satisfying correlation can be the fact that in the case of calorimetry the interaction parameter

is obtained from diluted solutions and in case of saturation solubility from saturated solutions.

Chapter6Polymers

63

Chapter 6:Polymers

Abstract

In this chapter is described a first series of experiments, which should show a transfer of the

principle of Self Incompatible Solvent from low molecular mass substances to high molecular

mass substances like polymers. Therefore a block copolymer: polybutadiene-block-poly(2-

vinylpyridine) is proposed as a model substance for a polymer which consists of two

incompatible parts. Anthracene is selected as a model substance for active ingredients.

Chapter6Polymers

64

6.1 Motivation

In the previous chapters a new theory Self Incompatible Solvent is thermodynamically

described and correlated to experimental data. The experimental data was obtained by using

precision solution calorimetry or saturation solubility. In both cases low molecular mass

substances were used and they showed good agreement between theoretical predictions and

the experimental data. Therefore, a further idea is to transfer this principle to high molecular

mass substances like polymers. According to the theory of Self Incompatible Solvent the

polymers should also consist of two different parts which “don’t like” each other. Such

polymers should appear as good solvents for various substances. However, the principle of a

Self Incompatible Solvent is based on the assumption that one may prevent the phase

separation by linking the two incompatible parts together by a covalent link. This assumption

is expected to fail if the two parts exceed a certain length or volume. Therefore, further

investigations were conducted in which two incompatible parts were linked together to form a

block copolymer and in which the length of the two blocks were varied systematically (while

the average composition was kept constant). A block copolymer composed of two extremely

short blocks is supposed to be homogeneous, increasing the chain length at some point will

induce phase separation. The expectation was that this phase separation would be

accompanied by a significant reduction in the ability of the block copolymer to dissolve an

additional ingredient.

Thus a block copolymer needs to be selected that fulfilled the requirements that it

should consist of two incompatible parts and that one should be able to tailor the length of the

two blocks. A crystalline substance was dissolved in this block copolymer - systematically

varying the block length - and its solubility was determined. The dependence of the solubility

on the block length was expected to look like depicted in the graph in Figure 32b. The

solubility should initially decreases weakly with increasing block length. Exceeding a

threshold of block length, phase separation of the block copolymer is expected and the

solubility should drop significantly. Following this step, the dependency of the solubility on

the block length again was expected to be only weak.

One task was to determine the saturation solubility of this crystalline substance in the block

copolymer. Because of the high viscosity of the block copolymer, (probably glassy state

occurs at room temperature), the separation of a saturated solution from an excess of non

crystallized solid was not feasible. One possible experimental procedure is to add an excess of

the substance to be dissolved, to wait for the equilibrium and to distinguish the dissolved

Chapter6Polymers

65

material from the undissolved one by spectroscopy. Thus, the solubility can be obtained from

a plot of the degree of crystallinity versus the concentration: The saturation solubility may be

obtained from the intercept of this line with the abscissa (see Figure 32b).

degree ofcrystallinity

concentrationsolubility

crystallinity


solubility

blocklength

solubility

length

a) b)



crystallinity




crystallinity


solubility

blocklength

solubility

length

solubility

blocklength

solubility

length

a) b)

Figure 32: Transfer of the principle of Self Incompatible Solvents to polymers: a) solubility versus block length of the copolymers b) degree of crystallinity versus concentration.

6.2 Polybutadiene-block-poly(2-vinylpyridine)

The idea of investigating the properties of phase separated block copolymers of

identical composition, but varying block length has already very nicely been pursued in a

completely different context by the group of Steven Holdcroft[23][24]. He investigated the ionic

conductivity of phase separated block copolymers composed of polystyrene and poly (sodium

styrene sulfonate) as function of the block length. The background of these investigations is

the expectation that ionic conductivity happens in water swollen regions composed of the poly

(styrene sulfonate) while the non water swollen styrene regions provide structural integrity

and limits the degree of swelling. Even at the identical overall compositions, the materials

composed of different block lengths yielded materials with variations in width and

connectivity of the conducting channels and thus showed systematic variations in overall

conductivity and the degree of swelling at a given humidity. These polymers that have been

prepared by Holdcroft's group were not exclusively block copolymers but to a large extend

graft copolymers that were prepared via stable free radical polymerization and emulsion

polymerization. The major focus in this research was not to investigate the transition between

homogeneous and phase separated state, but to influence the characteristic length scale of the

phase separated system by varying the block lengths. Thus, the preparation of extremely short

blocks was not in the focus of this research. It seemed not to be obvious that the synthetic

Chapter6Polymers

66

procedures used in his group would give a narrow chain length distribution and close to 90%

efficient coupling of both blocks even at comparatively low molar masses. Furthermore, in the

case of Holdcroft's polymers both compounds have a glass transition temperature exceeding

room temperature. In order to facilitate the formation of an equilibrium at least one of the

polymeric blocks should be non glassy at room temperature.

One polymerization technique, which makes a high degree of coupling possible, even

if the block lengths are short, is anionic polymerization. One especially favorable block

copolymer that can be synthesized with anionic

polymerization and fulfils the above mentioned

requirements is the block copolymer: polybutadiene-

block-poly(2-vinylpyridine) (see Figure 33). There is a

incompatibility between the two blocks, as shown

in eqn. 64:

-parameter = rep. ( polybutadiene - polyvinylpyridine)2

= (8.31[25]-10.70[26])2 = 5.71 (cal/cm3)1/2 eqn.64

This incompatibility can be increased further if the

pyridine block will be additionally quaternized. Polybutadiene is a polymer with a low glass

transition temperature of -15°C[27]. The block copolymer can be synthesized via anionic

polymerization with various block lengths by the subsequent addition of monomers and by the

variation of the ratio of monomer to initiator.

On the other hand anionic polymerization requires special equipment and experience;

therefore collaboration with Professor Axel Müller and the PhD Student Felix Schacher from

the Bayreuth University was initiated. The four polymers listed in Table 8 were synthesized

by Felix Schacher (with my assistance) in the laboratories of Prof. A. Müller in Bayreuth: The

following block lengths were intended: B10V5, B20V10, B30V15 and B70V35, where B

corresponds to polybutadiene block and V corresponds to poly(2-vinylpyridine). The results

obtained in the laboratory showed minor deviations form the intended block lengths, but all

have identical overall composition and a systematic variation of the block lengths was

achieved.

( )n( )m

N

( )n( )m

N

Figure 33: Polybutadiene-block-poly(2-vinylpyridine).

Chapter6Polymers

67

Table 8: Composition of synthesized polybutadiene-block-poly(2-vinylpyridine) polymers; B corresponds to polybutadiene block, V corresponds to poly(2-vinylpyridine), (Data received from Felix Schacher).Polymer B

[massfraction

%]

V[mass

fraction%]

DP of

B

DP of

V

Mw of

B

[g/mol]

Mw of

V

[g/mol]

Mw of

block

copolymer

[g/mol]

PDI Description

B10V5 52 48 14 7 800 750 1550 1.1 B14V71550

B20V10 45 55 24 15 1300 1600 2900 1.08 B24V152900

B30V15 50 50 30 15 1600 1600 3200 1.07 B30V153200

B70V35 46 54 73 26 3900 2700 6600 1.06 B73V266600

To check, whether phase separation between the blocks of the polymers occur, differential

scanning calorimetry (DSC) measurements were done. In the first series of experiments all

DSC measurements were done with a heating rate of 10 K/min. Due to the fact that after DSC

measurements with heating rate of 10 K/min it was still difficult to recognize whether in the

polymers B30V15 and B70V35 one or two glass transition temperatures occur, second set of

measurements was made with a higher heating rate of 20 K/min.

In Figure 34 the results of the DSC measurements for all synthesized polymers are

summarized. As can be seen, the polymers B30V15 and B70V35 have two glass transitions

temperatures (Tg). This indicates that in case of these two polymers phase separation occurs.

Nevertheless, the transition temperatures of poly(2-vinylpyridne) and polybutadiene in

polymers B30V15 and B70V35 are not identical with the values given in literature. In “Polymer

Handbook”[28][29][30][31] the value of the glass transition temperatures of poly(2-vinylpyridne)

and polybutadiene are 104°C and -15°C respectively. The prepared polymers B30V15 and

B70V35 show glass transition temperature for poly(2-vinylpyridne) at 25°C and at 50°C

respectively and for butadiene by -25°C. Prepared polymers have low molecular masses (see

Table 8) in comparison to the polymers given in “Polymer Handbook”. This can be the reason

for the deviation of the glass transition temperatures from literature and experiments.

Chapter6Polymers

68

Figure 34: Differential scanning calorimetry of the polymers shown in Table 8. Temperature program (with 10 K/min and 20 K/min):

1) Heating from -100 °C to 200 °C2) Hold 2 min by 200 °C

3) Cooling from 200 °C to -100 °C 4) Hold 2min by -100 °C

5) Heating from -100 °C to 200 °C ; only the last temperature ramp is shown.

In the next step, an active ingredient was chosen which shows two different

fluorescence signals, one for the dissolved form and the second for the crystalline form in

polymers. Fortunately, anthracene perfectly fits to these criteria. Anthracene is good soluble

in ethanol and the fluorescence spectrum of the dissolved anthracene is shown in Figure 35

(blue curve). It can be seen that the emission maximum of dissolved anthracene is at a

wavelength of 410 nm. Due to the fact that anthracene is almost insoluble in water (0.0001

g/L) it should be possible to measure the fluorescence of crystalline anthracene after

dispersing it in that solvent, see Figure 35 (black curve). Unfortunately, crystals of anthracene

are difficult to disperse in water, they cling to the wall of the corresponding containers and

have a tendency to cream and sediment, therefore it is not easy to prepare dispersions of an

exactly predetermined concentration. Another possibility is to measure the fluorescence of

crystalline antracene which is deposited onto glass plate, see Figure 35 (rose curve). The

green curve in Figure 35 shows the fluorescence spectra of anthracene in a mixture of

Chapter6Polymers

69

water/ethanol (20:1 by volume). The area in the violet frame in Figure 35 shows a linear

combination of spectra of dissolved and crystalline anthracene (55% dissolved, 45%

crystalline). This experiment shows that it is possible to detect crystalline and dissolved forms

anthracene via fluorescence spectroscopy.

Figure 35: Fluorescence spectras of anthracene.

In the next step it was necessary to detect the concentration of crystalline and

dissolved anthracene in the block copolymers which are described in this chapter. One way to

achieve this might be the preparation of several mixtures of the block copolymers and

anthracene, e.g. by co-dissolving both in a common volatile solvent and doctor blading this

mixture onto a suitable substrate and measuring the fluorescence spectra of each of these

samples. Due to the fact that anthracene would be dissolved in the mixture of solvent and the

polymer and not only in the pure polymer it is necessary to use another procedure for the

preparation of the anthracene solution in the polymer. More elegant might be to pursue a

method of high throughput screening following the procedures published by Professor

Amis[32][33]. First, a mixture of the dissolved block copolymer in THF and anthracene is doctor

bladed with a tilted doctor blade onto the glass plate to form a wedge. Then the dissolved

block copolymer in THF is doctor bladed onto the second glass plate again in the form of a

wedge. The thickness of the deposited film increases from 0 on the one side of the plate to

0.25 mm on the second side of the plate. Next, the solvent is evaporated from the both doctor

bladed solutions and after vitrification these both polymer wedges are bond together by

pressing the glass plates against each other in opposing orientation. The bonded polymer

wedges are then kept in an oven at 100°C for one week, to allow the anthracene to diffuse

400 450 500 550 6000

100200300400500600700800900

1000

Inte

nsity

[a.u

]

wavelength [nm]

anthracene in ethanol anthracene in water crystalline anthracene anthracene in

water/ethanol mixture (20:1 by volume)

400 450 500 550 6000

10

20

30

40

50

Inte

nsity

[a.u

]

wavelength [nm]

anthracene,

partially dissolved in

water/ethanol mixture (20:1)

linear combination

of spectra of dissolved and

crystalline anthracene

55% dissolved,

45% crystalline

Chapter6Polymers

70

from one wedge into the other and thus to form a vertically homogeneous layer that has a

lateral gradient in the anthracene concentration.

a.)

b.)

c.)

d.)

Figure 36: Procedure of the preparation of the block copolymer samples for the fluorescence: a) polymer wedge with anthracene and wedge with pure polymer b) wedges putted in opposing directions c) pressed glass plates with wedges d) bonded block copolymer wedges after one week at 100°C, verticallyhomogeneous layer.

To measure the fluorescence of dissolved and crystalline anthracene in the polymer, glass

plates need to be cut into stripes. This approach was pursued as sketched above.

Unfortunately, the desired advantage of high throughput screening did not pay of. To measure

the fluorescence using the available spectrometer, the glass plates needed top be cut into

stripes. During cutting, the glass plates broke in to irregular pieces that were too small to

mount them into the holder of the fluorescence spectrometer. Remaining bigger glass parts of

the plates moved against each other during the cutting process and this caused a smearing out

of the polymeric films and made the estimation of the concentrations unreliable.

At this point it was decided to stop working on this subproject and to transfer it into the hands

of my successor, Yana Wang.

a)

b)

c)

d)

Chapter7Experimental part

71

Chapter 7Experimental part

Abstract

This chapter contains experimental details for purifications and synthesis processes of all

model substances. Methods used for characterizations of all synthesized substances will be

shown. The Precision Solution Calorimetry measurements and the principle of this method

will be explained. Additionally the experimental data are summarized in table form.


72

7.1 Purification process of active ingredients

7.1.1 AnthraceneAnthracene was purchased by the Merck Company with a purity of 96% (weight percent).

Before use anthracene was purified by a chromatographic column. 200 ml Cyclohexane

(99 %, Sigma-Aldrich, used as received) are poured into a 500 ml round brown bottom flask,

which was connected to a dropping funnel with a pressure equalizing arm. The drooping

funnel was used as a chromatographic column filled with: glass wool, 30 g sea sand, 80 g

aluminium oxide 90 active neutral, 100 g anthracene and glass wool, see Figure 37 .

glass woolsea sand

aluminum oxide90 active neutral

glass wool

anthracene

cyclohexane

glass woolsea sand


glass wool

anthracene

cyclohexane

Figure 37: Schematic illustration of the used apparatus for purification of anthracene.

Cyclohexane is heated and reflux. The condensed solvent and flows with dissolved anthracene

through the chromatographic column. Afterwards the solution is cooled down to room

temperature and anthracene crystallize. The solvent is decanted and anthracene dried under

vacuum for 8 hours. Dried anthracene is pulverized and then for precision solution

calorimetry measurements.


73

7.1.2 9-anthracene carboxylic acid9-anthracene carboxylic acid (99 %, Merck) is recrystallized from ethanol (technical grade),

with a standard method described in literature[34]. In a 250 ml round bottom flask made from

brown glass 200 ml of ethanol and 150 g of 9-anthracene carboxylic acid were mixed. The

solution is refluxed for 6 hours. The solution was cooled down to room temperature and 9-

anthracene carboxylic acids crystallize. The solvent is decanted and the active ingredient is

dried under vacuum for 8 hours. Dried 9-anthracene carboxylic acid is pulverized and used for

precision solution measurements.

7.1.3 AcridineAcridine (96 %, Merck Company) was purified in two steps. In the first step, acridine is

recrystallized from n-heptane (99 %, Sigma-Aldrich, used as received). In a 250 ml round

bottom flask made from brown glass, 200 ml of n-heptane and 100 g of acridine are mixed

and stirred. The solution was refluxed for 6 hours and afterwards cooled to room temperature

and acridine crystallized. The solvent and impurities are decanted and the active ingredient

was dried under vacuum for 8 hours. In the second step, acridine was dissolved in a

water/ethanol mixture (ratio 1:1 by volume) and passed through column with activated

charcoal and aluminium oxide 90 active neutral, similar to the purification of anthracene

(Figure 38). Then acridine was recrystallized from the ethanol/water mixture.

glass woolsea sand


activated charcoal

glass wool

acridine

water/ethanol

glass woolsea sand


activated charcoal

glass wool

acridine

water/ethanol

Figure 38: Schematic illustration of the used apparatus for the purification of acridine.


74

150 ml of water and 150 ml of ethanol are poured into a 500 ml round brown bottom flask,

which was connected to a dropping funnel with a pressure equalizing arm. The drooping

funnel was used as a chromatographic column filled with glass wool, 30 g sea sand, 80 g

aluminum oxide 90 active neutral, 15 g activated charcoal (1-3 mm), 90 g acridine and glass

wool (Figure 38). The 1:1 mixture of water and ethanol refluxed. The condensed solvent and

flows with dissolved anthracene through the chromatographic column. Afterwards the

solution is cooled down to room temperature and anthracene crystallize. The solvent is

decanted and anthracene dried under vacuum for 8 hours. Dried anthracene is pulverized and

then for precision solution calorimetry measurements.

7.2 Used equipment

7.2.1 Nuclear Magnetic Resonance (NMR)

All spectra are obtained via a spectrometer from the Varion Company (type Oxford 400) at

room temperature in deuterium solvents.

o 1H NMR (CDCl3, 400 MHz, ppm) Solvent as internal standard:

CHCl3 = 7.26, Si(CH3)4 = 0.00

o 13C NMR (CDCl3, 100 MHz, ppm) Solvent as internal standard:

CHCl3 = 77.76, Si(CH3)4 = 0.00

7.2.2 Infrared Spectroscopy (IR)All spectra are obtained via a spectrometer from the Bruker Optics Company (type IFS-48).

All measurements are obtained at room temperature, on a KBr plates and in an atmosphere

almost free of water and carbon dioxide.

7.2.3 Fluorescence spectroscopyFor all fluorescence measurements the Luminescene Spectrometer LS 50B from the Perkin

Elmer Company is used. The solution with incomplete dissolved active ingredient was filtered

with a syringe filter (pore size 0.2 µm, PP, Whatman). After filtration, solutions are diluted in

ethanol. For every active ingredient standard concentrations are prepared and measured.

For anthracene the following parameters are used: Excitation: 360 nm; Begin: 370 nm; End:

620 nm; Speed: 240 nm/min.


75

As standards, solutions with the following concentrations of anthracene in ethanol are

prepared: 1.00 mg/l, 0.80 mg/l, 0.60 mg/l, 0.50 mg/l, 0.25 mg/l, 0.10mg/l. All spectra of the

standard solutions are shown in Figure 39 and the intensities of these concentrations

in Figure 40.

400 450 500 550 6000

100

200

300

400

500

600 Anthracene in ethanol

Inte

nsity

[a.u

]

wavelenght [nm]

1.00 mg/l 0.80 mg/l 0.60 mg/l 0.50 mg/l 0.25 mg/l 0.10 mg/l

Figure 39: Fluorescence spectra of the standard solutions of anthracene in ethanol.

0.0 0.2 0.4 0.6 0.8 1.00

100

200

300

400

500

600

Inte

nsity

[a.u

]

concentration [mg/l]

y = 542,87xR2 = 0,9986

Figure 40: Concentrations versus fluorescence intensity of the anthracene standard solutions.

From the calibration curve ( xy 87.542 , see Figure 40) the concentrations of the solutions

after Solution Calorimetry measurements are calculated. The correct values of H are

calculated with the equation given below:


76

lgclVgmmolkJHH

nce fluorescesubstanceliquid

e anthracenasurements SolCal me

// eqn. 64

Calibration curves for 9-anthracene carboxylic acid and acridine are prepared and afterwards

the concentrations are determined by the same method as for anthracene.

Parameters used for determination of the concentration of 9-anthracene carboxylic acid via

fluorescence are given below are given below:

Excitation: 360 nm; Begin: 370 nm; End: 620 nm; Speed: 240 nm/min

Parameters used for determination of the concentration of acridine via fluorescence are given

below:

Excitation: 340 nm; Begin: 350 nm; End: 620 nm; Speed: 240 nm/min

In all three cases the obtained concentrations are further used for the calculation of the heats

of dissolution from the corresponding solution calorimetry experiments and for the calculation

of the heats of dissolutions from the saturation solubility experiments.

7.2.4 Precision Solution Calorimetry

7.2.4.1 The principle of calorimetry

The reaction calorimetry is an often used method, which allows determination of the

quantitative change of enthalpy during chemical reactions or dissolution processes. In the

calorimeter, the heats of processes like dissolution are determined as a change of temperature

of the reaction vessel using the heat capacity of the filled vessel to compute the heat from the

temperature change. Therefore a comparison experiment, the so called calibration, is done

which allows to obtain the heat capacity of the calorimeter (Figure 41).


77

Temperature

Time

Temperature

Time

Figure 41: Calibration of a calorimetry experiment[35].

In the first step of this experiment only the temperature inside the calorimeter is detected (first

baseline). In the second step the calorimeter is electrically heated with known power

(calibration) and in the third step the temperature after heating is obtained (second baseline).

Although in the first and third step of the experiment no heat was produced, the baselines are

not horizontal, because the calorimeter is not completely insulated from the environment and

thus slowly approaches the same temperature as the surrounding. This approach is

proportional to the difference of the temperature between vessel and environment, with

involvement of this loss of heat.

7.2.4.2 SolCal Precision Solution Calorimeter

The enthalpies of mixing are measured in the SolCal Precision Solution Calorimeter

from the Thermometric Company (Sweden). This setup consist of the following main parts:

a) the calorimetric unit, b) calorimeter cylinder which is placed permanently in a TAM 2277

thermostat and holds the calorimetric unit when experiments are running, c) Solution

Calorimeter Module mounted in the 2280 TAM Accessory Interface, which is responsible for

the electronic functions, see Figure 42.


78

a) b) c)

Figure 42: Main part of SolCal Precision Solution Calorimeter a) calorimeter unitb) calorimeter cylinder c) Solution Calorimeter Module[35].

The calorimetric unit consists of two parts: i) stirrer system consisting of two further parts: a

stirrer motor and gold stirrer which hold glass ampoule (V = 1 ml) containing a sample, ii) a

reaction vessel (V = 25 ml) containing Peltier-elements which are permanently mounted in

two pockets descending from the top of the vessel (Figure 43).


79

Glas ampouleV = 1ml

Liquid

Glass vesselV = 25ml

Glas ampouleV = 1ml

LiquidGlas ampouleV = 1mlGlas ampouleV = 1ml

LiquidLiquid


Glass ampouleV= 1 ml Glass vessel

V= 25 ml

Liquid

Glas ampouleV = 1ml

Liquid


Glas ampouleV = 1ml

LiquidGlas ampouleV = 1mlGlas ampouleV = 1ml

LiquidLiquid


Glass ampouleV= 1 ml Glass vessel

V= 25 ml

Liquid

a) b)

Figure 43: Reaction vessel for SolCal Precision Solution Calorimetera) image[35] b) schematic illustration.

Before solution calorimetry measurements with active ingredients were done, a standard

experiment of dissolution of potassium chloride in water at 25 °C was made. From the

literature it is known that solHm (298.15 K) = (17.584 ± 0.017) kJ·mol–1[36][37][38].

0.0280 g of KCl is encased in the glass ampoule (volume of empty ampoule is 1 ml). The

ampoule is closed with the two-component epoxy glue (UHU plus directly solid). After 24

hours of glue solidification, the ampoule is mounted into the gold stirrer. Then the reaction

vessel is filled with 25 ml of distilled water. The gold stirrer with the mounted ampoule is

immersed into the water filled reaction vessel and they are screwed together. Afterwards the

glass vessel with the ampoule is attached to the gold stirrer and placed in the calorimetric

cylinder. The temperature profile of a calorimetric experiment is shown in Figure 44. The

measurement starts with a vessel temperature which is below the detection range of the

temperature sensor, line (A). The vessel is heated with a heat power of 500 mW to the

temperature that is detectable via the temperature sensor, line (B). The experiment consists of

the following parts: baseline (BL), first calibration (C), the reaction which is initiated by

breaking the ampoule (break), baseline (BL), second calibration (C) and baseline (BL), see

Figure 44. The final temperature at the complete equilibrium with the heat bath and the rate of

heat exchange with the bath are calculated by the SolCal program from the baseline. Both

calibrations should give comparable values.

The measurements with 0.0280 g of potassium chloride give a value of solHm = 17.536

kJ·mol–1.


80

Time [h]

Tem

pera

ture

[K]

Time [h]

Tem

pera

ture

[K]

Figure 44: Typical temperature profile of a calorimetric experiment; dissolution ofpotassium chloride in water at 25°C.

Parameters used for all break experiments:

One calibration before break; Calibration parameters:

Heat tPQ = 10 J

Power P = 500 mW

Time t = 20 s

Duration of baseline after calibration: 5 min

Break parameters

Duration: 5 min

Duration of baseline after break: 5 min

One calibration after break; Calibration parameters:

Heat tPQ = 10 J

Power P = 500 mW

Time t = 20 s

Duration of baseline after calibration: 5 min

Pause when temperature offset is greater than 500 mK

Stirring rate: 600 rpm

Heater power: 500 mW

Target temperature offset: 0 mK


81

7.3 Synthesis and purification of the model substances for compound 1 and 2

The number of schemes corresponds to the numbers in Table 3.

Scheme 1

Methyl valerate (99 %, Merck) and n-butyl acetate (99 %, Merck) were chromatographed on

an aluminumoxide 90 active basic (Merck) column before use for the solution calorimetry

experiment as a first pair (compound 1 and 2), see Figure 45.

a) b)

O

O

O

O

Figure 45: Structural formulas of a) methyl valerate b) n-butyl acetate.

Scheme 2Methyl pivalate (for synthesis, 98 % Merck) and tert-butylacetate (99 %, Aldrich) were

chromatographed on an aluminium oxide 90 active basic column (Merck) before use for the

solution calorimetry experiment as a second pair (compound 1 and 2), see Figure 46.

a) b)

O

O

O

O

Figure 46: Structural formula of a) methyl pivalate b) tert-butylacetate.

Scheme 3Methyl 3,3-dimethylbutonate (98 %, Alfa Aesar) and nonylacetate (97 %, Aldrich) were


solution calorimetry experiment as a third pair (compound 1 and 2), see Figure 47.

a) b)

O

O

O

O

Figure 47: Structural formula of a) methyl 3,3-dimethylbutonate b) nonylacetate.


82

Scheme 4Methyl decanoate (99 %, Aldrich) and 2,2-dimethyl-1-propylacetate (synthesized in the

laboratory) were chromatographed on an aluminium oxide 90 active basic column (Merck)

before use for the solution calorimetry experiment as a fourth pair (compound 1 and 2), see

Figure 48.

a) b)

O

O

O

O

Figure 48: Structural formula of a) methyl decanoate b) neopentyl acetate.

The synthesis procedure of neopentyl acetate is given below:

1 mol (88.15 g) neopentyl alcohol (99 %, Alfa Aesar) was mixed with 80 ml dry pyridine in a

500 ml three neck round bottom flask. The flask is immersed into an ice bath and connected

to a condenser and to a dropping funnel with a pressure equalizing arm. 0.9 mol (70.65 g,

64.05 ml) of cold acetyl chloride (99 %, for synthesis, Merck) was added from the dropping

funnel into the mixture of neopentyl alcohol and pyridine during 1hour. After addition of

acetyl chloride, the mixture was stirred for 1hour at 60 °C. After this time, the mixture was

poured into 300 g of ice. After phase separation; the upper layer (product) was washed with

250 ml of HCl (5 %) and with saturated potassium hydrogen carbonate solution. Afterwards it

was dried over magnesium sulfate (65.3 % yield). After purification neopentyl acetate is

characterized with IR- spectroscopy via a characteristic ester band (-COO-) at 1745.0 cm-1. In

Figure 49 is shown a 1H NMR spectrum and the following peaks are obtained, (CDCl3, 400

MHz, ppm): 3.76(s, 2H), 2.07(s, 3H), 0.94 (s, 9H). In Figure 50 is shown a 13C NMR

spectrum (CDCl3, 100 MHz, ppm): 171.20, 73.74, 31.14, 26.35, 20.85.


83

Figure 49: 1H NMR spectrum of neopentyl acetate.

Figure 50: 13C NMR spectrum of neopentyl acetate.

(ppm)255075100125150175200

171.

239

C

O

OCH31

5 23

4

1

2

3

4

5

4

4

(ppm)255075100125150175200

171.

239

C

O

OCH31

5 23

4

1

2

3

4

5

4

4

[ppm]

O

O

3 4

4

12

45

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

1.7 2.6 9.4

C

O

OCH3

1

2

312

3

3

3

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

1.7 2.6 9.4

C

O

OCH3

1

2

312

3

3

3O

O 3

3

3

12

[ppm]


84

Scheme 5Methyl trifluoroacetate (98%, for synthesis, Merck) and isopropyl acetate (99%, for synthesis,

Merck) were chromatographed on an aluminium oxide 90 active basic column (Merck) before

use for the solution calorimetry experiment as a fifth pair (compound 1 and 2), see Figure 51.

a) b)

O

O

F

F

F O

O

Figure 51: Structural formula of a) methyl trifluoroacetate b) isopropyl acetate.

Scheme 6Methyl trifluoroacetate (98%, for synthesis, Merck) and n-butyl acetate (99.5%, for analysis,

Merck) were chromatographed on an aluminium oxide 90 active basic column (Merck) before

use for the solution calorimetry experiment as a sixth pair (compound 1 and 2), see Figure 52.

a) b)

O

O

F

F

F O

O

Figure 52: Structural formula of a) methyl trifluoroacetate b) n-butyl acetate.

Scheme 7Methyl heptafluoroacetate (97 %, ABCR Fluorochemicals) and n-butyl acetate (99 %,

Aldrich) were chromatographed on an aluminium oxide 90 active basic column (Merck)

before use for the solution calorimetry experiment as a seventh pair (compound 1 and 2), see

Figure 53.

a) b)

O

O

FF

F FF

F O

O

Figure 53: Structural formula of a) methyl heptafluoroacetate b) n-butylacetate.

Scheme 8Methyl decanoate (98 %, Merck) and triethyleneglycol monomethylether acetate (synthesized

in the laboratory) were chromatographed on an aluminium oxide 90 active basic column


85

(Merck) before use for the solution calorimetry experiment as an eight pair (compound 1 and

2), see Figure 54.

a) b)

O

O

O

O

OO

O

Figure 54: Structural formula of a) methyl deaconate b) triethylene glycol monomethylether acetate.

The synthesis procedure of the triethylene glycol monomethylether acetate is given below.

2 mol (204.18 g) of freshly distilled acetanhydrid (98 %, Merck), 1 mol (164.20 g) anhydrous

triethylene glycol monomethylether (98 %, Merck) and 15 g Amberlite 200 (Merck) as

catalyst, were added into a 500 ml round bottom flask, which was connected to condenser.

The mixture was heated at 160 °C for 7 days. The product was distilled from the mixture

(50 °C, 1.4-1 mbar). The reaction yield was 78.3 %. The product was characterized with IR, 1H

NMR, and 13C NMR spectroscopy. The IR analysis showed characteristic ester band (-COO-)

at 1742.1 cm-1. In Figure 55 spectrum of the 1H NMR is shown: (CDCl3, 400 MHz, ppm);

4.24(t, 2H), 3.67(m, 8H), 3.55(t, 2H), 3.39(s, 3H), 2.08(s, 3H), and in Figure 56 spectrum of

the 13C NMR is shown: (CDCl3, 100 MHz, ppm): 171.00, 71.91, 70.59, 69.11, 63.61, 59.04,

20.97.

Figure 55: 1H NMR spectrum of triethylene glycol monomethylether acetate.

C

O

O

O

O

O

CH3

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.7 2.5 1.8 7.5 1.7

45

3

2

1

5

3

412 2

2 2

[ppm]

O

O

OO

O3

412

5 2

2

2


86

Figure 56: 13C NMR spectrum of triethylene glycol monomethylether acetate.

Scheme 9Methyl 3,3-dimethylbutonate (99 %, Alfa Aster) and triethyleneglycol monomethylether

acetate (synthesized in the laboratory, synthesis recipe, see scheme 8) were chromatographed

on an aluminium oxide 90 active basic column (Merck) before use for the solution

calorimetry experiment as a ninth pair (compound 1 and 2), see Figure 57.

a) b)

O

O

O

O

OO

O

Figure 57: Structural formula of a) methyl 3,3-dimethylbutonate b) triethylenglycol monomethylether acetate.

Scheme 10Methyl decanoate (98 %, Merck) and 2-(2-ethoxy ethoxy ethyl) acetate (98 %, Merck) were


solution calorimetry experiment as a tenth pair (compound 1 and 2), see Figure 58.

255075100125150175200

C

O

O

O

O

O

CH3

5

3

4

1

4

2

1

6

(ppm)

7

7

5

3

3 3

2 6O

O

OO

O5624

7 3

3

31

[ppm]


87

a) b)

O

O

O

O

OO

Figure 58: Structural formula of a) methyl decanoate b) 2-(2ethoxy ethoxy ethyl)acetate.

Scheme 11Methyl decanoate (98 %, Merck) and tetrahydrofurfuryl acetate (98 %, Alfa Aster) were


solution calorimetry experiment as an eleventh pair (compound 1 and 2), see Figure 59.

a) b)

O

O

O

O O

Figure 59: Structural formula of a) methyl decanoate b) tetrahydrofurfuryl acetate.

Scheme 12Methyl decanoate (98 %, Merck) and 1-(2 ethyl) pyrolidione acetate (synthesized in the


before use for the solution calorimetry experiment as a twelfth pair (compound 1 and 2), see

Figure 60.

a) b)

O

OO

O

N

O

Figure 60: Structural formula of a) methyl decanoate b)1-(2- ethyl) pyrolidione acetate.

The synthesis procedure of 1-(2- ethyl) pyrolidione acetate is given below:

2 mol (204.18 g) fresh distilled acetic anhydride (98 %, Merck), 1 mol (129.16 g) 1-(2-

hydroxyethyl-2- pyrolidione) (98%, Merck) and 15 g of Amberlite 200 (Merck) as catalyst

were mixed in 500 ml round bottom flask which was connected to condenser. The mixture

was heated at 160 °C for 7 days. The product was distilled from the mixture (100 °C, 3.2 -

1mbar) and the yield of the reaction was 70.6 %. IR spectroscopy showed a characteristic ester

bond (-COO-) at 1742.1 cm-1. The 1H NMR spectrum is shown in Figure 61: (CDCl3, 400


88

MHz, ppm): 4.21(t, 2H), 3.56(t, 2H), 3.48(t, 2H), 2.38(t, 2H), 2.05(m, 5H). In Figure 62 is

shown the 13C NMR spectrum: (CDCl3, 100 MHz, ppm): 174.80, 170.70, 61.73, 47.85,

41.61, 30.67, 20.83, 17.98.

Figure 61: 1H NMR spectrum of 1-(2- ethyl) pyrolidione acetate.

Figure 62: 13C NMR spectrum of 1-(2- ethyl) pyrolidione acetate.

255075100125150175200

C

O

OCH3 N

O

74

1

85

6

1 2

3

4

5

67

8

32

[ppm]

O

O

N

O

6

85

4

37

1

2

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

5.0 2.0 2.1 2.1 1.9

C

O

OCH3 N

O

14

5 5

4

1

2

32

3

5

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

5.0 2.0 2.1 2.1 1.9

C

O

OCH3 N

O

14

5 5

4

1

2

32

3

5

O

O

N

O

4

53

2

15

[ppm]


89

Scheme 13Methyl decanoate (98%, Merck) and 1,3- diethoxypropan-2-yl acetate (synthesized in the


before use for the solution calorimetry experiment as a thirteenth pair (compound 1 and 2),

see Figure 63.

a) b)

O

O

O

O

O

O

Figure 63: Structural formula of a) methyl deaconate b) 1,3- diethoxypropan-2-yl acetate

The synthesis of 1,3- diethoxypropan-2-yl acetate is given below:

2 mol (204.18 g) fresh distilled acetic anhydride (98 %, Merck), 1 mol (148.2 g)

1,3 diethoxypropanol (98 %, Merck) and 15 g Amberlite 200 (Merck) as catalyst, were mixed

in 500 ml round bottom flask, connected to an condenser. The mixture was heated at 160°C

for 7 days. The product was distilled from the mixture (42 °C, 1.4-1 mbar), and the reaction

yield was 75.4 %. IR spectrum showed characteristic ester bond (-COO-) at 1742.1 cm-1 In

Figure 64 is shown the 1H NMR spectrum: (CDCl3, 400 MHz, ppm): 5.12(m, 1H), 3.54(m,

8H), 2.10(s, 3H), 1.19 (t, 6H). In Figure 65 is shown the 13C NMR spectrum (CDCl3, 100

MHz, ppm): 170.60, 71.63, 68.96, 66.81, 21.25, 15.07.

Figure 64: 1H NMR spectrum of 1,3- diethoxypropan-2-yl acetate.

C

O

OO

OCH3

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

6.0 3.0 8.3 0.7

1

2

3

4

4

42

22

2

13

(ppm) [ppm]

O

O

O

O4

4

2

2

2

2

13


90

Figure 65: 13C NMR spectrum of 1,3- diethoxypropan-2-yl acetate.

Scheme 14N-butylacetate (99 %, Merck,) and benzylacetate (98%, Merck) were chromatographed on an

aluminium oxide 90 active basic column (Merck) before use for the solution calorimetry

experiment as a fourteenth pair (compound 1 and 2), see Figure 66.

a) b)

O

O

O

O

Figure 66: Structural formula of a) n-butylacetate b) benzyl acetate.

102030405060708090100110120130140150160170180190200

170.

628

C

O

OO

OCH3

4

5

3

32

2

1

51

6

3

4

4

66

[ppm]

O

O

O

O

6

6

4

4

3

3

25 1


91

7.4 Synthesis and purification of the model substances forSelf Incompatible Solvent

The number of schemes corresponds to the numbers in Table 3.

Scheme 1 – Self Incompatible SolventN-butyl valerate (98 %, Aldrich) was chromatographed on an aluminium oxide 90 active

basic column (Merck) before use for the solution calorimetry experiment as a first model

substance for Self Incompatible Solvent, see Figure 67.

O

O

Figure 67: Structural formula of n-butyl valerate.

Scheme 2 – Self Incompatible SolventTert-butyl pivalate (synthesized in the laboratory) was used for the solution calorimetry

experiment as a second model substance for Self Incompatible Solvent, see Figure 68.

O

O

Figure 68: Structural formula of tert-butyl pivalate.

The synthesis procedure of tert-butylpivalate is given below:

1.4 mol (103.8 g) tert-butyl alcohol (98 %, Merck) was mixed with 120 ml of dry pyridine in

a 500 ml three neck round bottom flask. The flask was immersed into an ice bad and

connected to a condenser and to a dropping funnel with pressure equalizing arm. 1 mol

(120.58 g, 123 ml) of cold pivaloyl chloride (98 %, Merck) was added from the dropping

funnel into the mixture during 1 hour. After addition of pivaloyl chloride, the mixture was

stirred for 1 hour at 80 °C. After this time, the mixture was poured into 500 g of ice. After

addition into ice, phase separation occurred, where upper layer was the product. The product

was washed with 250 ml of 5 % HCl and saturated potassium hydrogen carbonate solution,

and dried over magnesium sulfate (81.3 % yield). IR spectrum showed the characteristic ester

bond (-COO-) at 1727.7 cm-1. The 1H NMR spectrum is shown in Figure 69: (CDCl3, 400

MHz, ppm): 1.43(s, 3H), 1.15(s, 3H). In Figure 70 the 13C NMR spectrum is shown:

(CDCl3, 100 MHz, ppm): 177.90, 79.43, 39.12, 27.89, 27.11.


92

Figure 69: 1H NMR spectrum of tert-butyl pivalate.

Figure 70: 13C NMR spectrum of tert-butyl pivalate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

3.0 3.1

C

O

O

12

1

1

1

2

2

2

[ppm]

2

O

O

2

2

1

1

1

255075100125150175200

C

O

O1

2

1

4

3

2

5

3

4

4

4

5

5

5

[ppm]

3

4

5

O

O

5

5

4

41 2


93

Scheme 3 – Self Incompatible SolventNonyl 3,3-dimethylbutanoate (synthesized in the laboratory) was used for the solution

calorimetry experiment as a third model substance for Self Incompatible Solvent,

see Figure 71.

O

O

Figure 71: Structural formula of nonyl 3,3-dimethylbutanoate.

The synthesis procedure of nonyl 3,3-dimethylbutanoate is given below:

1.5 mol (216.38 g) of 1-nonanol (99%, Fluka) was mixed with 200 ml of dry pyridine in a

1000 ml three neck round bottom flask. The flask was immersed into ice bath and connected

to a condenser and to a dropping funnel with pressure equalizing arm. 1.4 mol (191.24 g,

197.4 ml) of cold 3,3-dimethylbutanoyl chloride (99%, Merck) was added from the dropping

funnel into the mixture during 1 hour. After addition of 3,3-dimethylbutanoyl chloride, the

mixture was stirred 1 hour at 60 °C. After this time, the mixture was poured into 800 g of ice

where phase separation occurred; upper phase was the product. The product was washed with

250 ml of 5 % HCl and saturated potassium hydrogen carbonate solution, and dried over

magnesium sulfate (81.3 % yield). IR spectrum showed the characteristic ester bond (-COO-)

at 1730.6 cm-1. The 1H NMR spectrum is shown in Figure 72; (CDCl3, 400 MHz, ppm):

4.05(t, 2H), 2.19(s, 2H), 1.62(m, 2H), 1.27(m, 12H), 1.03(m, 9H), 0.88(t, 3H). In Figure 73 is

shown the 13C NMR spectrum; (CDCl3, 100 MHz, ppm): 172.40, 64.11, 48.04, 31.81,

29.60, 29.18, 28.64, 25.96, 22.62, 14.05.


94

Figure 72: 1H NMR spectrum of nonyl 3,3-dimethylbutanoate.

Figure 73: 13C NMR spectrum of nonyl 3,3-dimethylbutanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.6 9.4 12.2 2.2 1.9 1.7

5

34

1

2

61

C

O

O 3

4 4 4

4 4 4

62

5

5

5

[ppm]

4 4O

O4

4 4433 6

2

55

5

255075100125150175200

1

C

O

O

3 4

5

8

6 6 9

2

9

8

8

67

10

1

8

3

4

2

6

7 6 5 10

[ppm]

3O

O2 1

04

88

8

17 56

6 966


95

Scheme 4 – Self Incompatible SolventNeopentyl decanoate (synthesized in the laboratory) was used for the solution calorimetry

experiment as a fourth model substance for Self Incompatible Solvent, see Figure 74.

O

O

Figure 74: Structural formula of neopentyl decanoate.

The synthesis procedure of neopentyl decanoate is given below:

1.2 mol (105.78 g) neopentyl alcohol (99%, Alfa Aestar) was mixed with 120 ml of dry

pyridine in 500 ml three neck round bottom flask. The flask was putted into an ice bad and

connected to a condenser and to a dropping funnel with pressure equalizing arm. 1 mol

(190.72 g, 205.1 ml) of cold decanoyl chloride (99 %, Aldrich) was added from the dropping

funnel into the mixture during 1 hour. After addition of decanoyl chloride, the mixture was

stirred for 1 hour at 60 °C. After this time the mixture is poured into 300 g of ice and then

phase separation occurred; upper layer was the product. The product was washed with 250 ml

of 5 % HCl, saturated potassium hydrogen carbonate solution and dried over magnesium

sulfate (81.3 % yield). IR spectrum showed characteristic ester band (-COO-) at 1739.2 cm-1.

In Figure 75 the 1H NMR spectrum is shown; (CDCl3, 400 MHz, ppm): 3.77(s, 2H), 2.32(t,

2H), 1.64(m, 2H), 1.27(m, 12H), 0.94(s, 9H), 0.879(m, 3H). In Figure 76 the 13C NMR

spectrum is shown; (CDCl3, 100 MHz, ppm): 173.90, 73.50, 34.38, 31.81, 31.22, 29.38,

26.39, 25.02, 22.62, 14.04.


96

Figure 75: 1H NMR spectrum of neopentyl decanoate.

Figure 76: 13C NMR spectrum of neopentyl decanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.7 9.3 12.4 1.6 1.6 1.8

1

3

42

5

6

C

O

O

5

5

5

12

3444

4 4 46

[ppm]

4

4O

O4 4

4 43

36

15

5

5

255075100125150175200

C

O

O

5

72

1

2

3

71

4 36

46 8

9

10

7

78

55

5 59

10

[ppm]

4 O

O

5 5 310 1

7

7

7

9 5 5 8 2 6


97

Scheme 5 – Self Incompatible SolventIsopropyl trifluoroacetate (98 %, Fluorochem) was chromatographed on an aluminium oxide

90 active basic column (Merck) before use for the solution calorimetry experiment as a fifth

model substance for Self Incompatible Solvent, see Figure 77.

O

O

F

F

F

Figure 77: Structural formula of isopropyl trifluoroacetate.

Scheme 6 – Self Incompatible SolventN-butyl trifluoroacetate (97%, ABCR) was chromatographed on an aluminium oxide 90

active basic column (Merck) before use for the solution calorimetry experiment as a sixth


O

O

F

F

F

Figure 78: Structural formula of n-butyl trifluoroacetate.

Scheme 7 – Self Incompatible SolventN-butylheptafluorobutanoate (97%, ABCR) was chromatographed on an aluminium oxide 90

active basic column (Merck) before use for the solution calorimetry experiment as a seventh


O

O

FF

F FF

F

Figure 79: Structural formula of n-butylheptafluorobutanoate.


98

Scheme 8 – Self Incompatible SolventTriethylene glycol monomethyl decanoate (synthesized in the laboratory) was used for the

solution calorimetry experiment as an eighth model substance for Self Incompatible Solvent,

see Figure 80.

O

O

OO

O

Figure 80: Structural formula of triethylene glycol monomethyl decanoate.

The synthesis procedure of triethylene glycol monomethyl decanoate is given below:

1 mol (172.27 g) decanoic acid (99 %, Alfa Aesar), 2 mol (328.4 g) triethylene glycol

monomethyl ether for synthesis (97 %, Merck), 300 ml toluol (98 %, Merck) and 15 g

Amberlite 200 (Merck) as catalyst were mixed in 1000 ml two neck round bottom flask. The

flask was connected to a water separator[39] and to a condenser. The mixture was heated to

160 °C till 18 ml of water was obtained (14 days). After this time Amberlite 200 were filtered

and the product was distilled from the mixture (100 °C, 1.3-1 mbar). The yield of the reaction

was 89.2 %. IR spectrum shown characteristic ester bond (-COO-) at 1736.3 cm-1. In Figure

81 the 1H NMR spectrum is shown; (CDCl3, 400 MHz, ppm): 4.23(t, 2H), 3.67(m, 8H),

3.55(m, 2H), 3.38(s, 3H), 2.33(t, 2H), 1.62(m, 2H), 1.26(m, 12H), 0.88(t, 3H), In Figure 82

the 13C NMR spectrum is shown; (CDCl3, 100 MHz, ppm): 173.60, 71.75, 70.40, 69.03,

63.18, 58.85, 34.02, 31.70, 29.25, 24.74, 22.50, 13.94.

Figure 81: 1H NMR spectrum of triethylene glycol monomethyl decanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.7 12.4 1.8 1.7 2.5 1.9 8.5 1.7

C

O

O

O

O

O

2

1

4

35

4

5

6

78

8

7 7 7

7 7 7

6 3

1

2 2

2 2O

O

OO

O7 7 7 6 1

27 7 7 58 2

2 2

3 4

[ppm]


99

Figure 82: 13C NMR spectrum of triethylene glycol monomethyl decanoate.

Scheme 9 – Self Incompatible SolventTriethylene glycol monomethyl tertbutanoate (synthesized in the laboratory) was used for the

solution calorimetry experiment as a ninth model substance for Self Incompatible Solvent, see

Figure 83.

O

O

OO

O

Figure 83: Structural formula of triethylene glycol monomethyl tertbutanoate.

The synthesis procedure of triethylene glycol monomethyl tertbutanoate is given below:

1.6 mol (185.86 g) 3,3 dimethylbutanoic acid (98 %, Alfa Aesar), 1.3 mol (213.46 g) of

triethylene glycol monomethyl ether (98 %, for synthesis, Merck), 300 ml of toluol (98 %,

Merck) and 15 g of Amberlite 200 (Merck) as catalyst were mixed in 1000 ml two neck round

bottom flask. The flask was connected to a water separator and to a condenser. The mixture

was heated to 160 °C till 18 ml of water was obtained in the water separator (14 days). After

this time Amberlite 200 were filtered and the product was distilled from the mixture (112 °C,

2.4-1 mbar) (90.4% yield). IR spectrum showed characteristic ester bond (-COO-) at

1733.4cm-1. In Figure 84 the 1H NMR spectrum is shown; (CDCl3, 400 MHz, ppm): 4.17(t,

2H), 3.60(m, 8H), 3.50(m, 2H), 3.32(s, 3H), 2.16(s, 2H), 0.97(s, 9H). In Figure 85 the 13C

255075100125150175200

C

O

O

O

O

O

4

58

7

12

10119

7

6

3

1

2

1

5 3

10

9

9

9

9

8

11

12

4 3

3

2 6

[ppm]

O

O

OO

O11

9 9 10

5

38 9 9 712

4

3 3

2 6


100

NMR spectrum is shown; (CDCl3, 100 MHz, ppm): 172.2, 77.40, 71.92, 70.61, 69.21,

62.98, 59.03, 47.84, 30.72, 29.60.

Figure 84: 1H NMR spectrum of triethylene glycol monomethyl tertbutanoate.

Figure 85: 13C NMR spectrum of triethylene glycol monomethyl tertbutanoate.

C

O

O

O

O

O

4

6

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

9.0 1.7 2.5 1.7 8.1 1.9

3

2

1

4

5

6

156

2 2

23 26

[ppm]

5

16

O

O

OO

O

6

6 3 2

2 2

2 4

102030405060708090100110120130140150160170180190200

172.

238

C

O

O

O

O

O

(ppm)

3

2

1

1

45 6

78

94

8

5

9

10

972

6 3

9

104

[ppm]

O

O

OO

O6

93 4

89

9 5 4 2 710 1


101

Scheme 10 – Self Incompatible SolventDiethylene glycol monoethyl decanoate (synthesized in the laboratory) was used for the

solution calorimetry experiment as a tenth model substance for Self Incompatible Solvent, see

Figure 86.

O

O

OO

Figure 86: Structural formula of diethylene glycol monoethyl decanoate.

The synthesis procedure of diethylene glycol monoethyl decanoateis given below:

1 mol (172.27 g) of decanoic acid (99 %, Alfa Aesar), 1.75 mol (234.815 g) diethylene glycol

monoethylether (98 %, for synthesis, Merck,), 300 ml toluol (98 %, Merck) and 15 g

Amberlite 200 (Merck) as catalyst were mixed in a 1000 ml two neck round bottom flask. The

flask was connected to a water separator and to a condenser. The mixture was heated to

160 °C till 18 ml of water was obtained in the water separator (14 days). After this time

Amberlite 200 were filtered and the product was distilled from the mixture (temperature 110

°C, pressure 2.4-1 mbar) (90.4 % yield). IR spectrum showed characteristic ester bond

(-COO-) at 1736.3 cm-1. In Figure 87 the 1H NMR spectrum is shown; (CDCl3, 400 MHz,

ppm): 4.19(m, 2H), 3.59 (m, 6H), 3.48 (t, 2H), 2.28(t, 2H), 1.58(t, 2H), 1.22(m, 15H), 0.83(t,

3H). In Figure 88 the 13C NMR spectrum is shown; (CDCl3, 100 MHz, ppm): 173.70, 70.51,

69.69, 69.09, 66.55, 63.21, 34.08, 31.74, 29.15, 24.78, 22.54, 15.03, 13.97.


102

Figure 87: 1H NMR spectrum of diethylene glycol monoethyl decanoate.

Figure 88: 13C NMR spectrum of diethylene glycol monoethyl decanoate.

C

O

O

O

O

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.8 15.2 1.8 1.9 5.9 1.9 1.9

1

5

43

27

6

1

2 2

2 3

6

4

5

666

666

7

[ppm]

O

O

OO

5 1

4

6

27

6 6

6 6 6 22 3

6

C

O

O

O

O

255075100125150175200

1

543

2 6

7

8

9

10

11

12

13

1

5

4 2

3 6

1210

799

99

8

11

13

[ppm]

O

O

OO

10 5

7

11

413

9 9

8 9 9 2

3 6

121


103

Scheme 11 – Self Incompatible SolventTetrahydrofurfuryl decanoate (synthesized in the laboratory) was used for the solution

calorimetry experiment as an eleventh model substance for Self Incompatible Solvent, see

Figure 89.

O

O O

Figure 89: Structural formula of tetrahydrofurfuryl decanoate.

The synthesis procedure of tetrahydrofurfuryl decanoate is given below:

1 mol (172.27 g) of decanoic acid (99 %, Alfa Aesar), 2 mol (196.20 g) of tetrahydrofurfuryl

alcohol (99 %, Aldrich), 300 ml of toluene (98 %, Merck) and 15 g of Amberlite 200 (Merck)

as catalyst were mixed in 1000 ml two neck round bottom flask. The flask was connected to a

water separator and to a condenser. The mixture was heated to 160 °C till 18 ml of water was

obtained in the water separator (14 days). After this time the Amberlite 200 were filtered and

the product was distilled from the mixture (118 °C, 2.4-1 mbar) (90.4 % yield). IR-spectrum

shown characteristic ester bond (-COO-) at 1742.1 cm-1. In Figure 90 the 1H NMR spectrum

is shown (CDCl3, 400 MHz, ppm): 4.11(m, 2H), 3.97(m, 1H), 3.80(m, 2H), 2.30(t, 2H),

1.88(m, 2H), 1.59(m, 4H), 1.22(m, 12H), 0.84(t, 3H). In Figure 91 the 13C NMR spectrum is

shown; (CDCl3, 100 MHz, ppm): 173.80, 76.48, 68.38, 66.24, 34.14, 31.79, 29.19, 27.92,

25.60, 24.87, 22.60, 14.04.


104

Figure 90: 1H NMR spectrum of tetrahydrofurfuryl decanoate.

Figure 91: 13C NMR spectrum of tetrahydrofurfuryl decanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.9 12.5 3.5 2.2 2.0 1.6 1.0 1.8

C

O

O

O

32

1

1

4

56

7

8

7

48 7

57

7 7

72

6

6

3

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.9 12.5 3.5 2.2 2.0 1.6 1.0 1.8

C

O

O

O

32

1

1

4

56

7

8

7

48 7

57

7 7

72

6

6

3

[ppm]

O

O O7 1

78

7

7 6

3

2 64

7

7

5

255075100125150175200

173.

809

(ppm)

32

1

4 5

6 11

12

9

C

O

O

O

111 7 72

8

3

9

10

7810

4

12 6 57 7

255075100125150175200

173.

809

(ppm)

32

1

4 5

6 11

12

9

C

O

O

O

111 7 72

8

3

9

10

7810

4

12 6 57 7

[ppm]

O

O O7 4

612

11

7 8

3

2 105

7

7

91


105

Scheme 12 – Self Incompatible Solvent2-pyrrolidinone decanoate (synthesized in the laboratory) was used for the solution

calorimetry experiment as a twelfth model substance for Self Incompatible Solvent, see

Figure 92.

O

O

N

O

Figure 92: Structural formula of 2-pyrrolidinone decanoate.

The synthesis procedure of 2-pyrrolidinone decanoate is given below:

1 mol (172.27 g) of decanoic acid (99 %, Alfa Aesar), 2 mol (258.32 g) of 1-(2-

hydroxyethyl)-2-pyrrolidinone (98 %, Aldrich), 300 ml of toluene (98 %, Merck) and 15 g

Amberlite 200 (Merck) as catalyst were mixed in 1000 ml two neck round bottom flask. The

flask was connected to a water separator and to a condenser. The mixture was heated to

160 °C till 18 ml of water were obtained in the water separator (14 days). After this time

Amberlite 200 were filtered and the product was distilled from the mixture

(130 °C, 8.0-2 mbar) (90.4 % yield). IR spectrum showed characteristic ester bond

(-COO-) at 1736.3 cm-1. In Figure 93 the 1H NMR spectrum is shown; (CDCl3, 400 MHz,

ppm): 4.17(t, 2H), 3.50(t, 2H), 3.42(t, 2H), 2.34(t, 2H), 2.26(t, 2H), 1.99(m, 2H), 1.57(m,

2H), 1.22(m, 12H), 0.84(t, 3H). In Figure 94 the 13C NMR spectrum is shown; (CDCl3, 100

MHz, ppm): 175.20, 173.50, 61.55, 47.87, 41.69, 34.16, 30.71, 29.20, 24.81, 22.61, 18.04,

14.06.


106

Figure 93: 1H NMR spectrum of 2-pyrrolidinone decanoate.

Figure 94: 13C NMR spectrum of 2-pyrrolidinone decanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

2.6 12.4 1.7 1.6 1.9 1.7 1.8 2.0 1.8

C

O

O N

O

1 54

32

67

8

9

1

2 36

5

4

7

888

888

9

[ppm]

O

O

N

O

18

9 8

8

8

8

8

7

4 2

36

5

255075100125150175200

C

O

O N

O

12

3

65

47

1113

12108

13

9

21

3

4 512

6

7

10

9

8

8

8

8

11

[ppm]

O

O

N

O

3111

38

8

8

8

9

10 7 4

51261

1


107

Scheme 13 – Self Incompatible Solvent1,3-diethoxy-2-yl decanoate (synthesized in the laboratory) was used for the solution

calorimetry experiment as a thirteenth model substance for Self Incompatible Solvent, see

Figure 95.

O

O

O

O

Figure 95: Structural formula of 1,3-diethoxy-2-yl decanoate.

The synthesis procedure of 1,3-diethoxy-2-yl decanoateis given below:

1 mol (172.27 g) of decanoic acid (99 %, Alfa Aesar), 2 mol (296.4 g) of 1,3-diethoxy-2-

propanol (96 %, Merck), 300 ml of toluene (98 %, Merck) and 15 g of Amberlite 200 (Merck)

as catalyst were mixed in a 1000 ml two neck round bottom flask. The flask was connected to

a water separator and to a condenser. The mixture was heated to 60 °C till 18 ml of water was

obtained in the water separator (14 days). After this time the Amberlite 200 were filtered and

the product was distilled from the mixture (115 °C, 1.3-1 mbar) (89.4 % yield). IR spectrum

showed the characteristic ester bond (-COO-) at 1736.3cm-1. In Figure 96 the 1H NMR

spectrum is shown; (CDCl3, 400 MHz, ppm): 5.13(m, 1H), 3.56(m, 8H), 2.34(t, 2H),

1.63(m, 2H), 1.26(m, 12H), 1.18(t, 6H), 0.87(t, 3H). In Figure 97 the 13C NMR spectrum is

shown; (CDCl3, 100 MHz, ppm): 173.30, 71.24, 69.01, 66.71, 34.40, 31.81, 29.37, 24.97,

22.61, 15.03, 14.05.


108

Figure 96: 1H NMR spectrum of 1,3-diethoxy-2-yl decanoate.

Figure 97: 13C NMR spectrum of 1,3-diethoxy-2-yl decanoate.

0.51.01.52.02.53.03.54.04.55.05.56.06.57.07.58.0

0.7 8.0 1.7 1.8 12.5 6.3 2.7

(ppm)

4

5

6

31

2

26

7

C

O

OO

O1

57

4

3 2

2 265

55

55

[ppm]

O

O

O

O5

75

5

4

355

5 12

22

2

6

6

255075100125150175200

1

3

C

O

OO

O

1

5

7

43

2

28

6

8

7 5

9

9

10

11

3

4

4

1010

7

7 7

611

[ppm]

O

O

O

O9

11

7

7

8

567

7 1 23

34

4

10

10


109

Schema 14 – Self Incompatible SolventBenzyl pentanoate (98%, Alfa Aesar) was chromatographed on an aluminium oxide 90 active

basic column (Merck) before use for the solution calorimetry experiment as a fourteenth


O

O

Figure 98: Structural formula of benzyl pentanoate.

7.5 Results from SolCal CalorimeterAll experimental data and results are summarized in Table 9, Table 10 and Table 11. To

obtain the heat of mixing of compound 1 in compound 2, 0.4 ml of compound 1 was enclosed

in glass ampoules and after the break of ampoule compound 1 was mixed with 25 ml of

compound 2. Because of this in the following tables in the columns “compound 1 in

compound 2” are given only the values for Q[J]. These values are used for further

calculations. Sometimes, the amount of active ingredient that was enclosed in the ampoule

was not completely dissolved in the liquid compound. Solubility of completely dissolved

compounds during SolCal experiment is needed to get the correct values of H1 [kJ/mol]. For

these samples, values of H1 [kJ/mol] are calculated from fluorescence measurements (see

chapter 7.2.3 and appendix 1). The values of H1 that are calculated from the fluorescence are

labeled with “*”.

Table 9: Measured values with anthracene in the SolCol Calorimeter.Scheme

Nr.

anthracene in

compound 1

anthracene in

compound 2

anthracene

in 50/50

mixtures

anthracene

in

SIS

compound 1

in

compound 2

m1 [g] 0.0342 0.0360 0.0343 0.0270 -

H1 [kJ/mol] 23.192 22.978 24.131 14.138 -

Q1 [J] 4.450 4.641 4.644 2.142 -0.214

m2 [g] 0.0255 0.0261 0.0269 0.0241 -

H2 [kJ/mol] 22.422 21.329 21.508 16.384 -

1.

Q2 [J] 3.208 3.123 3.246 2.215 -0.179

2. m1 [g] 0.0368 0.0360 0.0341 0.0119 -


110

H1 [kJ/mol] 24.481 24.581 23.283 16.121 -

Q1 [J] 5.055 4.965 4.454 1.076 -0.106

m2 [g] 0.0306 0.0253 0.0267 0.0267 -

H2 [kJ/mol] 20.908 20.573 21.199 15.656 -

Q2 [J] 3.589 2.920 3.176 2.345 -0.132

m1 [g] 0.0328 0.0316 0.0375 0.0088 -

H1 [kJ/mol] 23.604 18.386 18.631 11.805 -

Q1 [J] 4.344 3.261 3.920 0.583 0.659

m2 [g] 0.0290 0.0286 0.0273 0.0274 -

H2 [kJ/mol] 21.300 17.072 18.841 14.249 -

3.

Q2 [J] 3.466 2.739 2.886 2.190 0.721

m1 [g] 0.0321 0.0303 0.0350 0.0303 -

H1 [kJ/mol] 18.902 22.199 20.030 15.591 -

Q1 [J] 3.404 3.774 3.933 2.650 2.015

m2 [g] 0.0305 0.0263 0.0278 0.0269 -

H2 [kJ/mol] 16.070 19.922 19.830 15.632 -

4.

Q2 [J] 2.750 2.940 3.662 2.359 2.102

m1 [g] 0.0368 0.0329 0.0321 0.0209 -

H1 [kJ/mol] 44.740* 24.825 32.657 20.545 -

Q1 [J] 5.801 4.586 5.881 2.409 -0.201

m2 [g] 0.0372 0.0293 0.0319 0.0207 -

H2 [kJ/mol] 45.105* 24.020 33.194 20.588 -

5.

Q2 [J] 5.921 3.949 5.978 2.391 -0.105

m1 [g] 0.0368 0.0360 0.0307 0.0099 -

H1 [kJ/mol] 44.740* 22.978 30.715 16.537 -

Q1 [J] 5.801 4.641 5.290 0.919 0.2009

m2 [g] 0.0372 0.0261 0.0312 0.0097 -

H2 [kJ/mol] 45.105* 21.329 32.657 16.694 -

6.

Q2 [J] 5.921 3.123 5.310 0.909 0.217

m1 [g] 0.0266 0.0360 0.0337 0.0040 -

H1 [kJ/mol] 27.480* 22.978 24.152 11.318 -

7.

Q1 [J] 1.734 4.641 5.420 0.254 0.227


111

m2 [g] 0.0272 0.0261 0.0039 0.0042 -

H2 [kJ/mol] 28.010* 21.329 22.806 10.143 -

Q2 [J] 1.820 3.123 4.990 0.239 0.235

m1 [g] 0.0321 0.0306 0.0309 0.0314 -

H1 [kJ/mol] 18.902 18.986 21.003 9.638 -

Q1 [J] 3.404 3.260 3.641 1.698 10.341

m2 [g] 0.0305 0.0330 0.0302 0.0115 -

H2 [kJ/mol] 16.070 19.010 19.267 9.861 -

8.

Q2 [J] 2.750 3.421 3.612 0.636 10.414

m1 [g] 0.0328 0.0306 0.0309 0.0256 -

H1 [kJ/mol] 23.604 18.986 21.003 6.049 -

Q1 [J] 4.344 3.260 3.641 0.869 6.053

m2 [g] 0.0290 0.0330 0.0302 0.0234 -

H2 [kJ/mol] 21.300 19.010 19.267 6.366 -

9.

Q2 [J] 3.466 3.421 3.612 0.836 6.124

m1 [g] 0.0321 0.0365 0.0311 0.0293 -

H1 [kJ/mol] 18.902 20.821 19.011 16.199 -

Q1 [J] 3.404 4.264 3.317 2.663 7.807

m2 [g] 0.0305 0.0274 0.0281 0.0289 -

H2 [kJ/mol] 16.070 21.024 17.370 15.807 -

10.

Q2 [J] 2.750 3.232 2.738 2.563 7.519

m1 [g] 0.0321 0.0318 0.0310 0.0489 -

H1 [kJ/mol] 18.902 20.880 19.140 8.252 -

Q1 [J] 3.404 3.725 3.329 2.263 7.685

m2 [g] 0.0305 0.0259 0.0310 0.0424 -

H2 [kJ/mol] 16.070 19.429 18.216 8.565 -

11.

Q2 [J] 2.750 2.812 3.168 2.364 7.550

m1 [g] 0.0321 0.0304 0.0350 0.0424 -

H1 [kJ/mol] 18.902 18.708 20.030 4.279 -

Q1 [J] 3.404 3.191 3.933 1.018 11.439

m2 [g] 0.0305 0.0304 0.0350 0.0396 -

12.

H2 [kJ/mol] 16.070 18.815 19.843 4.510 -


112

Q2 [J] 2.750 3.209 3.896 1.002 12.390

m1 [g] 0.0321 0.0377 0.0338 0.0441 -

H1 [kJ/mol] 18.902 18.730 19.095 5.121 -

Q1 [J] 3.404 3.962 3.621 1.267 4.702

m2 [g] 0.0305 0.0376 0.0337 0.0412 -

H2 [kJ/mol] 16.070 18.811 19.005 5.170 -

13.

Q2 [J] 2.750 3.979 3.604 1.195 4.613

m1 [g] 0.0342 0.0351 0.0342 0.0634 -

H1 [kJ/mol] 23.192 22.626 23.325 10.408 -

Q1 [J] 4.450 4.456 4.475 3.702 0.723

m2 [g] 0.0255 0.0277 0.0281 0.0563 -

H2 [kJ/mol] 22.422 23.107 22.606 10.580 -

14.

Q2 [J] 3.208 3.591 3.564 3.342 0.627

Table 10: Measured values with acridine in the SolCol Calorimeter.Scheme

Nr.

acridine in

compound 1

acridine in

compound 2

acridine in

50/50

mixtures

acridine in

SIS

compound 1

in

compound 2

m1 [g] 0.1024 0.1161 0.1058 0.0899 -

H1 [kJ/mol] 18.194 16.288 17.038 16.453 -

Q1 [J] 10.397 10.553 10.059 8.254 -0.214

m2 [g] 0.1124 0.0982 0.1073 0.0932 -

H2 [kJ/mol] 18.798 16.763 16.888 16.111 -

1.

Q2 [J] 11.791 9.186 10.112 8.379 -0.179

m1 [g] 0.1076 0.1346 0.1050 0.0989 -

H1 [kJ/mol] 18.075 18.043 18.177 18.686 -

Q1 [J] 10.853 13.553 10.651 10.313 -0.106

m2 [g] 0.1235 0.1152 0.1032 0.0954 -

H2 [kJ/mol] 17.587 18.852 17.946 18.370 -

2.

Q2 [J] 12.121 12.011 10.335 10.125 -0.132

m1 [g] 0.1064 0.1088 0.0578 0.1254 -

H1 [kJ/mol] 16.664 14.983 16.050 8.285 -

3.

Q1 [J] 9.894 9.097 5.179 4.863 0.659


113

m2 [g] 0.1123 0.1172 0.1090 0.1438 -

H2 [kJ/mol] 16.134 14.979 15.900 8.093 -

Q2 [J] 10.111 9,797 9.671 6.494 0.721

m1 [g] 0.1151 0.1174 0.1103 0.1196 -

H1 [kJ/mol] 15.169 16.817 16.869 7.286 -

Q1 [J] 9.743 10.017 10.383 4.863 2.015

m2 [g] 0.1281 0.1121 0.1087 0.1093 -

H2 [kJ/mol] 14.775 16.012 17.021 7.763 -

4.

Q2 [J] 10.565 10.017 10.325 4.735 2.102

m1 [g] 0.1068 0.1161 0.1072 0.1312 -

H1 [kJ/mol] 20.749 16.288 20.900 16.999 -

Q1 [J] 12.366 10.553 12.503 12.446 -0.201

m2 [g] 0.1132 0.0982 0.1034 0.1299 -

H2 [kJ/mol] 20.186 16.763 20.868 16.552 -

5.

Q2 [J] 12.752 9.186 12.041 11.998 -0.105

m1 [g] 0.1068 0.1139 0.1258 0.1382 -

H1 [kJ/mol] 20.749 17.580 21.382 16.737 -

Q1 [J] 12.366 11.714 15.011 12.908 0.2009

m2 [g] 0.1132 0.1334 0.1222 0.1353 -

H2 [kJ/mol] 20.186 17.067 21.502 16.867 -

6.

Q2 [J] 12.752 12.705 14.663 12.735 0.217

m1 [g] 0.1065 0.1161 0.1222 0.103 -

H1 [kJ/mol] 17.663 16.288 20.446 15.321 -

Q1 [J] 10.497 10.553 13.942 8.806 0.227

m2 [g] 0.0987 0.0982 0.1242 0.1154 -

H2 [kJ/mol] 17.978 16.763 20.289 15.061 -

7.

Q2 [J] 9.902 9.186 14.062 9.699 0.235

m1 [g] 0.1151 0.1253 0.1061 0.1155 -

H1 [kJ/mol] 15.169 14.446 14.101 7.875 -

Q1 [J] 9.743 10.101 8.349 5.076 10.341

m2 [g] 0.1032 0.199 0.1045 0.1169 -

8.

H2 [kJ/mol] 15.630 14.813 14.080 7.872 -


114

Q2 [J] 9.001 9.911 8.211 5.135 10.414

m1 [g] 0.1064 0.1253 0.1145 0.0965 -

H1 [kJ/mol] 16.663 14.446 15.037 11.828 -

Q1 [J] 9.894 10.101 9.608 6.370 6.053

m2 [g] 0.1029 0.199 0.1137 0.0934 -

H2 [kJ/mol] 16.940 14.813 14.872 11.993 -

9.

Q2 [J] 9.727 9.911 9.436 6.251 6.124

m1 [g] 0.1151 0.1098 0.1077 0.1198 -

H1 [kJ/mol] 15.169 15.012 14.225 8.214 -

Q1 [J] 9.743 9.198 8.549 5.491 7.807

m2 [g] 0.1072 0.1037 0.1059 0.1065 -

H2 [kJ/mol] 15.464 14.812 14.099 8.955 -

10.

Q2 [J] 9.251 8.572 8.332 5.322 7.519

m1 [g] 0.1151 0.1089 0.0578 0.094 -

H1 [kJ/mol] 15.169 14.673 13.255 8.367 -

Q1 [J] 9.743 8.917 4.275 4.389 7.685

m2 [g] 0.1072 0.1089 0.1325 0.087 -

H2 [kJ/mol] 15.464 14.615 14.252 8.850 -

11.

Q2 [J] 9.251 8.882 10.538 4.297 7.550

m1 [g] 0.1151 0.1013 0.1101 0.0607 -

H1 [kJ/mol] 15.169 14.042 13.580 8.674 -

Q1 [J] 9.743 7.938 8.343 2.938 11.439

m2 [g] 0.1072 0.1011 0.1079 0.0594 -

H2 [kJ/mol] 15.464 14.162 14.059 9.047 -

12.

Q2 [J] 9.251 7.990 8.465 2.999 12.390

m1 [g] 0.1151 0.1099 0.0636 0.2004 -

H1 [kJ/mol] 15.169 14.828 13.238 6.498 -

Q1 [J] 9.743 9.094 4.698 7.267 4.702

m2 [g] 0.1072 0.1174 0.1037 0.1885 -

H2 [kJ/mol] 15.464 14.603 13.339 6.667 -

13.

Q2 [J] 9.251 9.567 7.719 7.024 4.613

14. m1 [g] 0.1132 0.1161 0.0646 0.0608 -


115

H1 [kJ/mol] 14.239 16.288 16.567 13.200 -

Q1 [J] 8.994 10.553 5.972 4.486 0.723

m2 [g] 0.1058 0.0982 0.1141 0.0591 -

H2 [kJ/mol] 14.590 16.763 16.120 13.108 -

Q2 [J] 8.614 9.186 10.264 4.323 0.627

Table 11: Measured values with 9 – anthracene carboxylic acid in the SolCol Calorimeter.Scheme

Nr.

9-anthracene

carboxylic

acid in

compound 1

9-anthracene

carboxylic

acid in

compound 2

9-anthracene

carboxylic

acid in

50/50

mixtures

9-anthracene

carboxylic

acid in

SIS

compound 1

in

compound 2

m1 [g] 0.0718 0.0613 0.0611 0.0989 -

H1 [kJ/mol] 15.059 15.993 13.405 13.495 -

Q1 [J] 4.865 4.411 3.685 6.005 -0.214

m2 [g] 0.0673 0.0625 0.0602 0.0972 -

H2 [kJ/mol] 15.758 16.112 13.294 13.703 -

1.

Q2 [J] 4.772 4.531 3.601 5.993 -0.179

m1 [g] 0.0668 0.0636 0.0616 0.1007 -

H1 [kJ/mol] 16.273 13.985 12.157 9.629 -

Q1 [J] 4.891 4.002 3.370 4.363 -0.106

m2 [g] 0.0701 0.0615 0.0622 0.0983 -

H2 [kJ/mol] 15.646 14.401 12.105 9.534 -

2.

Q2 [J] 4.935 3.985 3.388 4.217 -0.132

m1 [g] 0.0576 0.0606 0.0705 0.0700 -

H1 [kJ/mol] 17.792 10.702 13.497 8.099 -

Q1 [J] 4.611 2.918 4.281 2.551 0.659

m2 [g] 0.0521 0.0613 0.0688 0.065 -

H2 [kJ/mol] 17.921 10.873 13.500 8.018 -

3.

Q2 [J] 4.201 2.999 4.179 2.345 0.721

m1 [g] 0.0600 0.066 0.0690 0.038 -

H1 [kJ/mol] 11.664 14.426 12.496 6.492 -

4.

Q1 [J] 3.149 4.284 3.880 1.110 2.015


116

m2 [g] 0.0632 0.0645 0.0653 0.035 -

H2 [kJ/mol] 11.731 14.386 12.059 6.344 -

Q2 [J] 3.336 4.175 3.543 0.999 2.102

m1 [g] 0.0122 0.0613 0.0657 0.00425 -

H1 [kJ/mol] 47.382* 15.993 19.061 24.578 -

Q1 [J] 2.601 4.411 5.635 0.0470 -0.201

m2 [g] 0.0110 0.0602 0.0615 0.00399 -

H2 [kJ/mol] 48.935* 15.886 19.305 25.121 -

5.

Q2 [J] 2.422 4.303 5.342 0.451 -0.105

m1 [g] 0.0122 0.0648 0.648 0.0143 -

H1 [kJ/mol] 47.382* 16.706 19.794 24.962 -

Q1 [J] 2.601 4.871 5.771 1.614 0.2009

m2 [g] 0.0110 0.0635 0.652 0.0129 -

H2 [kJ/mol] 48.935* 16.286 20.077 25.463 -

6.

Q2 [J] 2.422 4.653 5.890 1.478 0.217

m1 [g] - - - - -

H1 [kJ/mol] - - - - -

Q1 [J] - - - - 0.227

m2 [g] - - - - -

H2 [kJ/mol] - - - - -

7.

Q2 [J] - - - - 0.235

m1 [g] 0.0600 0.0652 0.0699 0.0793 -

H1 [kJ/mol] 11.664 10.421 10.905 3.663 -

Q1 [J] 3.149 3.057 3.430 1.307 10.341

m2 [g] 0.0632 0.0649 0.0643 0.0686 -

H2 [kJ/mol] 11.731 10.277 11.095 3.638 -

8.

Q2 [J] 3.336 3.001 3.210 1.123 10.414

m1 [g] 0.0576 0.0652 0.0647 0.1023 -

H1 [kJ/mol] 17.792 10.421 12.464 12.748 -

Q1 [J] 4.611 3.057 3.629 5.868 6.053

m2 [g] 0.0521 0.0649 0.0613 0.1011 -

H2 [kJ/mol] 17.921 10.277 12.001 12.190 -

9.

Q2 [J] 4.201 3.001 3.310 5.545 6.124


117

m1 [g] 0.0600 0.0655 0.0642 0.0878 -

H1 [kJ/mol] 11.664 10.746 9.900 10.621 -

Q1 [J] 3.149 3.167 2.860 4.196 7.807

m2 [g] 0.0632 0.0623 0.0626 0.0792 -

H2 [kJ/mol] 11.731 10.866 9.657 11.129 -

10.

Q2 [J] 3.336 3.046 2.720 3.966 7.519

m1 [g] 0.0600 0.0732 0.0671 0.0568 -

H1 [kJ/mol] 11.664 10.545 8.642 6.761 -

Q1 [J] 3.149 3.473 2.609 1.728 7.685

m2 [g] 0.0632 0.0676 0.0633 0.0511 -

H2 [kJ/mol] 11.731 10.064 8.388 7.146 -

11.

Q2 [J] 3.336 3.061 2.389 1.643 7.550

m1 [g] 0.0600 0.065 0.0669 0.0566 -

H1 [kJ/mol] 11.664 1.935 1.127 -1.279 -

Q1 [J] 3.149 0.566 0.339 -0.326 11.439

m2 [g] 0.0632 0.0615 0.0639 0.0532 -

H2 [kJ/mol] 11.731 1.796 0.997 -1.301 -

12.

Q2 [J] 3.336 0.497 0.278 -0.311 12.390

m1 [g] 0.0600 0.0645 0.0716 0.0515 -

H1 [kJ/mol] 11.664 12.636 11.386 7.203 -

Q1 [J] 3.149 3.667 3.668 1.669 4.702

m2 [g] 0.0632 0.0627 0.0693 0.0499 -

H2 [kJ/mol] 11.731 12.208 11.392 7.117 -

13.

Q2 [J] 3.336 3.444 3.552 1.598 4.613

m1 [g] 0.0633 0.0613 0.0639 0.0344 -

H1 [kJ/mol] 19.040 15.993 16.144 8.806 -

Q1 [J] 5.423 4.411 4.642 1.363 0.723

m2 [g] 0.0597 0.0602 0.0627 0.0365 -

H2 [kJ/mol] 19.697 15.886 15.781 9.158 -

14.

Q2 [J] 5.291 4.303 4.452 1.504 0.627


118

7.6 PolymersAll Anionic polymerizations[40][41] were done by Felix Schacher in the laboratories of

Prof. Axel Müller at the University of Bayreuth. (Experimental details and characterization

by NMR, Size exclusion chromatography and Maldi-ToF, courtesy of Felix Schacher and

Prof. Axel Müller)

The polymerizations were prepared under nitrogen atmosphere in a thermostated glass

reactor (Büchi). The equipment for the polymerization is shown in Figure 99. The reactor

with 600 ml of tetrahydrofurane (THF) was cooled to -70 °C and the burette for the first

monomer butadiene to -20 °C. Sec-butyl lithium in hexane (sec-BuLi, Aldrich) which was

used as an initiator was added at this low temperature to the THF. Then the condensed

butadiene was added fast. After the addition of the monomer the temperature was slowly

increased to -15 °C. During this time the conversion of the butadiene (disappearance of the

double bounds) was observed via FT-NIR. When the double bonds of the butadiene were

disappeared a sample (so called precursor) was taken from the reactor. The GPC analysis of

the precursor gives information about the polybutadiene block. The reactor was cooled to

-70 °C and 2-vinylpyridine (Aldrich) (purification was performed with triethyl aluminum

solution in 1.0 M in hexane from Aldrich and subsequent condensation on a high vacuum

line) was added via a syringe. The color of the reaction mixture changed from yellow to deep

red. After the addition the reaction was kept at -70 °C for 1 hour. After this time, the living

chain ends were terminated with 5 ml of degassed isopropyl alcohol which was added via a

syringe. The color of the reaction mixture changed from deep red to colorless. The polymer

was purified by precipitation in a solution of 0.1 mol/L NaOH.


119

Figure 99: Equipment for the anionic polymerisation under FT-NIR monitoring of the monomer conversion.

To obtain 20 g of B20V10 the following amounts of the substances were used: 10 g

butadiene (15.38 ml, 185 mmol), 9.7 g 2-vinylpyridine (10 ml, 92.50 mmol), 6.60 ml sec-

BuLi 1.4 M in hexane (9.25 mmol).

The synthesized polymers and the results from GPC (polydispersity, PDI), MALDI –

ToF (molecular mass, Mn) and NMR (ratio of butadiene to 2-vinylpyridine, B:V) are

summarized in Table 8.

Vacuum

N2

Butadiene

Reactor forbutadienepurification

Condensed

Butadiene

Solventdistillation

setup

THF THF

FT-NIR


120

B10V5

Figure 100: GPC results from the polybutadiene-block-poly(2-vinylpyridine)B14V7

1550 sample.

Figure 101: MALDI - ToF results from the polybutadiene-block-poly(2-vinylpyridine)B14V7

1550 sample.

0 2000 4000 60000

5000

10000

15000

20000

25000

30000

35000

40000

B14V71550a.i.

1000 points smoothed of B14V7

1550 a.i

Det

ecto

r cou

nts

[a.u

.]

m/z

33 34 35 36 37 38 39 40

Elution volume [mL]

B precursor

B14V71550


121

Figure 102: 1H NMR results from the polybutadiene-block-poly(2-vinylpyridine)B14V7

1550 sample.

B20V10


2900 sample.

1234567891011121314

0.1 1.0

[ppm]

32 33 34 35 36 37 38 39

B precursor

B24V152900

Elution volume [mL]


122


2900 sample.


2900 sample.

1234567891011121314

1.0 0.4

[ppm]

0 2000 4000 60000

5000

10000

15000

20000

Det

ecto

r cou

nts

[a.u

.]

m/z

B24

V15

2900a.i. 1000 points smoothed

of B24V152900 a.i


123

B30V15


3200 sample.


3200 sample.

0 2000 4000 60000

2000

4000

6000

8000

10000

12000

14000

16000 B30V153200a.i.

1000 point smoothed of B30V15

3200 a.i

Det

ecto

r cou

nts

[a.u

.]

m/z

30 31 32 33 34 35 36 37 38

Elution volume [mL]

B precursor

B30V153200


124


3200 sample.

B70V35

26 28 30 32 34 36 38

Elution volume [mL]

B precursor B73V26

6600

26 28 30 32 34 36 38

Elution volume [mL]

B precursor B73V26

6600


6600 sample.

1234567891011121314

1.0 0.2

[ppm]


125

2000 4000 6000 8000 100000

1000

2000

3000

4000

5000

Det

ecto

r cou

nts

[a.u

.]

m/z

B73

V26

6600a.i. 1000 points smoothed

of B73V266600 a.i

Figure 110: MALDI-ToF results from the polybutadiene-block-poly(2-vinylpyridine)B73V26

6600 sample.


6600 sample.

1234567891011121314

1.0 0.3

[ppm]

Appendix 1

126

Appendix 1

Calculation of k12

Scheme 12: methyl deaconate as compound1 and 1-(2 ethylpyrolidione acetate) as compound

2 it is chosen as a example:

2112 V

Hk mix

JHmix 439.11

36373621

372

361

104.251041025104

1025

mmmVVVmV

mV

0160984011

984010542

1052

12

35

351

1

..

.m.

m.VV

336

2112 0629055

016098401042543911 m/kJ.

..m.J.

VHk mix

All calculations for k12 are done by analogy to the example showed above.

Calculation of ksis3, k13, k23 ki3

An example of measurements, where anthracene (compound 3) is dissolved in butyl valerate

(Scheme 1 for Self Incompatible Solvent, SIS) is chosen:

3

33

sis

fussolsis V

HHk

36

3)(3

)(3

)(3

1025

/1000000

/9.28

027.0142.2

mV

mg

molkJHgm

JH

sis

anthracene

anthracenefus

anthracene

sol

Appendix 1

127

0011.01

9989.0105027.2

105.2

105027.21000000

027.01025

38.4/24.178

027.0/9.28

3

35

35

35336)(3

)(3

)(3)(3

sis

sissis

anthracenesis

anthracene

anthracenefusanthracenefus

mm

VV

mmmVVV

Jmolg

gmolkJM

mHH

3

353

33

/03.828970011.09989.0105027.2

38.4142.2

mkJm

JJV

HHk

sis

fussolsis

All calculations for ksis3, k13, k23 are done by analogy to the example showed above.

Values for 3Hfus for other active ingredients:

molkJH acridinefus /58.18)(3[42]

molkJH acidcarboxylicanthracenefus /14.39)9(3[43]

Calculation of Hmix by usage of dissolved amounts of compound 3obtained from fluorescence measurements

From the calibration curve (see Figure 40):

xy 87.542 . ionconcentratIntensity 87.542

The concentrations of the solutions after SolCal measurements are calculated from the

following equations:

87.542Intensityionconcentrat dilution factor

69.30Intensity

dilution factor =10000

lglmglmgc /565.0/33.5651000087.542

/69.30

molkJH entsl measuremfrom SolCa /427.1gm e anthracen 0733.0lV cesliquid sub 025.0tan

lgc escencefrom fluor /565.0The value of solution enthalpy is calculated from the following equation:

molkJlgl

gmolkJlgclV

gmmolkJHH

escencefrom fluorcesliquid sub

e anthracenentsl measuremfrom SolCasol /41.7

/565.0025.00733.0/427.1

//

tan

Appendix 2

128

Appendix 2

Change in chemical potential upon dissolution

- binary mixture, both molecules have identical size, both compounds are liquids:

3333 lnln NVkRTNRTNG iiiisol

23

23

33

23

3

33

3 )()()(

i

i

i

i

NNN

NNNNN

NNN

dNd

dNd

23

23

3

3

3

32

3 )()(0

i

i

ii NNN

NNN

NNN

dNd

dNd

233

233

33

3333

3

23

33

23

233

333

3

ln

ln

1ln

11ln

)(

)(1ln0

)(1ln

ii

iiii

iiiiii

iiiiii

i

i

iii

i

i

iii

i

isolsol

VkRT

VkRT

VkRTRTRT

VkRTRTRT

NNN

NVk

NNN

RTNRT

NNN

RTNRTdN

Gd

323

33 )1(

ln0 isol kRT

- binary mixture, both molecules differ in size, compounds one is crystalline as pure

substance

3333333 lnln NGVNkRTNRTNG fusiiiisol

23

32

33

333333

33

33

33

3 )()()(

i

ii

ii

ii

ii NNVNV

VNVNVVNVNVNV

VNVNVN

dNd

dNd

Appendix 2

129

23

32

33

33

33

3

33 )()(0

i

ii

ii

i

ii

ii NN

VNVVNVNVNV

VNVNVN

dNd

dNd

32

333

3

33333

3

333333

33

3

3233

33333

233

3

233

3

333

33

1ln

1ln

1ln

)(

)(1ln0

)(1ln

GVkVVRTRT

GVkVVRTRTRT

GVVkVV

RTRTRT

GVNVN

VVNVNVk

VNVNVNV

RTNRT

VNVNVNV

RTNRTdN

Gd

fusiii

i

fusiii

ii

fusiiii

ii

fusii

iiii

ii

ii

iii

ii

iisolsol

3233

33

33

3 )1(

1)1(ln0 i

fusi

sol kV

GVVRT

Chapter 8Conclusion

130

Chapter 8:Conclusion

The theory describing the effect of Self Incompatible Solvent was derived. As

described in chapter 2, this theory assumes that, if a solvent is composed of two incompatible

parts and a third substance is dissolved in this solvent, then there is a negative contribution to

the enthalpy of mixing. According to this theory, this negative contribution should be the

higher, the stronger the unfavorable interactions between these two parts of the solvents are.

In chapter 2, theoretical predictions for Self Incompatible Solvent were described. These are

based on mean field theory, which assumes that all the interactions of a molecule with its

environment are given by the composition weighted mean value of the corresponding

interactions with the compounds comprising the environment. Additionally, the theory

assumes that the interactions are symmetric, that pair-wise interactions are assumed to be

proportional to the product of a given property, the interaction parameter of each partner of

the pair. A qualitative description of the mixing phenomena based on the mean field theory

was also shown. In two thermodynamic cycles, predictions of the thermodynamics of mixing

for binary and ternary mixtures were shown, taking into account the side constraints that (i)

all compounds are liquids and (ii) one compound is crystalline.

To prove this theory, two models systems are proposed in chapter 3:

1.) Physical 50/50 mixture - a system which was as close as possible to the

theoretical prediction given in the theory. In this model system the proposed

examples for compound 1 and 2 were homogenously, physically mixed and

Chapter 8Conclusion

131

then a crystalline compound 3 was added into this preexisting mixture. The

enthalpy of mixing of compound 3 with this mixture was measured.

2.) Self Incompatible Solvents – in this system compound 1 and 2 were forced

to exist together in one molecule via covalent bound. Therefore the enthalpy

of mixing of compound 3 with one compound – the Self Incompatible

Solvent – was measured. In this thesis, the Self Incompatible Solvent

consists of an alcohol group of the first compound which is linked together

via esterification with carboxylic acid of the compound two.

All interaction parameters of these two models described in chapter 3, were calculated from

the enthalpy of mixing obtained from the precision solution calorimetry.

On the basis of these two model systems it was shown that the thermodynamic

predictions, which were described in chapter 2 showed a good agreement with the

experiments. The results obtained via the precision solution calorimetry proved the theory and

it was concluded that unfavorable interactions between two mixed compounds are favorable

for the dissolution of a third one. That means that the self incompatibility of the solvents

improves the dissolution process.

It was shown that the model physical 50/50 mixture agreed much better with

theoretical prediction in comparison to the model of the Self Incompatible Solvent. The

covalent bond which connects these two incompatible parts of the Self Incompatible Solvent

was an additional modification of the system; this modification was not taken into account in

the theory, therefore this deviation was acceptable.

In chapter 4 was described a dissolution experiment which is based on Hess’s law.

This experiment allows the usage of the standard equipment of the precision solution

calorimetry for any desired mixing ratio of a two compounds. The expectation that the

character of the chosen solvents for this experiment had no significant effect on the results

was confirmed. Additionally, comparison of the results obtained from the solution calorimetry

and the theoretically calculated results based on theories of Van Krevelen and Hoy, was

made. The correlation of the results obtained from experimental and the theoretical data was

not satisfying. This is in accordance with an attempt to correlate theoretical predictions

calculated from the two independent methods of i) van Krevelen and ii) Hoy. The predictions

of these theoretical systems deviated from each other to approximately the same extend as

they deviated from the experimental data that were collected in the course of this thesis (see

section 1.5). Therefore the dissatisfying correlation of the results obtained from the

experimental and theoretical data was not surprising.

Chapter 8Conclusion

132

In chapter 5 was shown that the saturation solubility can be used to obtain the

interactions parameters instead of the precision solution calorimetry: all substances which

were used for Self Incompatible Solvents were saturated with a crystalline compound 3

decanted from the insoluble residue. Finally the concentration of the saturated solutions was

evaluated using the fluorescence spectroscopy. These results were used for further calculation

of interaction parameters. According to the theory, the interaction parameters obtained from

the precision solution calorimetry and saturation solubility should have the same values. The

comparison of the results from these two methods showed not the same correlation.

Nevertheless, even if the saturation solubility is used, the principle of Self Incompatible

Solvent seems to work, the self incompatibility of the solvents favoured the solubility of the

active ingredients, see chapter 5, section 5.4.

In the chapter 6 the principle of Self Incompatible Solvent is transferred from low

molecular mass substances to high molecular mass substances like polymers. The idea behind

the Self Incompatible Solvent is that a covalent link between the incompatible parts should

prevent phase separation. However, this only holds true, if the incompatible parts are not too

large. Above a certain length or volume of the incompatible parts micro phase separation has

to be expected and this should drastically diminish the expected favourable effect. To check

this hypothesis a series of block polymers of polybutadiene-block-poly(2-vinylpyridine),

which have two incompatible parts were polymerised, systematically varying the block

length of the two blocks. A first set of experiments which should show a dependence of the

solubility from the block length were designed.

This work showed that the principle of the Self Incompatible Solvents works. The

theory was proved by three different series of experiments: i) by using solution calorimetry

and the calculation of the interaction parameters between compounds 3 and preexisting

mixture of compound 1 and 2, ii) by using solution calorimetry and the calculation of the

interaction parameters between compounds 3 and the Self Incompatible Solvents, iii) from the

saturation solubility experiments and calculation of the interaction parameters between

compounds 3 and the Self Incompatible Solvents. These all results concluded in ones: that the

self incompatibility of the solvents improves the dissolution process.

133

References:[1] Hamley, I.; Introduction to Soft Matter, John Wiley & Sons, Chichester, 2000,Chapter 3.14.[2] Bergenståhl B. Surface Chemistry in Food and Feed. In Handbook of Applied Surfaceand Colloid Chemistry, Holmberg, K.; Shah, D. O.; Schwuger. M. J.;John Wiley & Sons,Chichester, 2002, Vol. 1, Chapter 2.

[3] Holmberg K. Surface Chemistry in Paints. In Handbook of Applied Surface and ColloidChemistry, Holmberg, K.; Shah, D. O.; Schwuger. M. J.;John Wiley & Sons, Chichester,2002, Vol. 1, Chapter 6.[4] Malmsten, M. Surface Chemistry in Pharmacy. In Handbook of Applied Surface andColloid Chemistry, Holmberg, K.; Shah, D. O.; Schwuger. M. J.;John Wiley & Sons,Chichester, 2002, Vol. 1, Chapter 1.

[5] Tadros, T. F. Surface Chemistry in Agriculture. In Handbook of Applied Surface andColloid Chemistry, Holmberg, K.; Shah, D. O.; Schwuger. M. J.;John Wiley & Sons,Chichester, 2002, Vol. 1, Chapter 4.[6] Tiberg, F.; Daicic J.; Fröberg J. Surface Chemistry of Paper. In Handbook of AppliedSurface and Colloid Chemistry, Holmberg, K.; Shah, D. O.; Schwuger. M. J.;John Wiley& Sons, Chichester, 2002, Vol. 1, Chapter 7.[7] Atkins, P.; de Paula, J. Physical Chemistry, Oxford University Press, Oxford, 2006,Chapter 5.[8] Meyer, D. Surfaces, Interfaces, and Colloids: Principles and Application, SecondEdition, John Wiley & Sons, New York, 1999, Chapter 4.[9] Hamley, I. Introduction to Soft Matter, John Wiley & Sons, Chichester, 2000,Chapter 3.[10] Atkins, P.; de Paula, J. Physical Chemistry, Oxford University Press, Oxford, 2006,Chapter 18.

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134

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[17] van Krevelen, D. W. Properties of Polymers Their Correlation with ChemicalStructure; Their Numerical Estimation And Prediction From Additive GroupContributions, Elsevier Science B. V., Amsterdam, 1990, Chapter 7.

[18] Rabek, J. F. Experimental methods in polymer chemistry, John Wiley & Sons,Chichester, 1980, Chapter 2, pp. 38.

[19] Elias, H. G. Polymere;: von Monomeren und Makromolekülen zu Werkstoffen; eineEinführung, Hüthing & Wepf, Heidelberg, 1996, Chapter 5, pp. 176.

[20] Dobias, B. “Coagulation and Flocculation; theory and applications”, Marcel Dekker,New York, 1993, pp. 169.

[21] Armarego, W. L. F., Perrin, D. D., Purification of laboratory chemicals, 1996,Butterworth-Heinemann-Press, Oxford, pp. 72.

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[23] Ding, J.; Chuy, C.; Holdcroft, S. Chem. Mater. 2001, 13, 2231.

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[27 ]Houston, D. J.; Ali, S. A. M. J. Appl. Polym. Sci. 1994, 52, 1129.

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AcknowledgmentsIt is a pleasure to thank the many people who made this thesis feasible, although I known thata few lines are too short to make a complete description of my deep gratitude.It is difficult to overstate my gratitude to my Ph.D. supervisor Prof. Werner A. Goedel. I wishto thank him for who he is as a professor. His belief and sincerity, his efforts in understandinga student’s personality, his enthusiasm, his encouragement, his enormous knowledge and hisgreat efforts to explain things clearly and simply, taken together, make him a great mentor. Hehelped to make chemistry fun for me.I also wish to express my appreciation to Dr. Sebastian Koltzenburg (BASF SE), forproviding many ideas and helpful discussions and suggestions. I would like to express mythanks to Dr. Ingrid Martin (BASF SE) for valuable discussions at the beginning of my thesis.I wish to express my cordial appreciation to Prof. Axel Müller and Dr. Felix Schacher(University of Bayreuth) for synthesis of polymers via anionic polymerization and theircharacterization by NMR, Size Exclusion Chromatography and Maldi-ToF.My sincere thanks are due to Prof. Stefan Spange, who allows me to use the great centrifugein his department. My special thanks go to a Ph.D. student Susanne Anders in Prof. StefanSpange group, for all DSC measurements.I am delighted to express my gratitude for NMR measurements by Michael Jahr, JensSeltmann and Robert Ullmann.A very sincere appreciation goes to the students: Tina Seifert and Arnd Nehrkorn for help bySolution Calorimetry measurements and by synthesis.The last years, it was a great and wonderful time for me, which I had spend in warm and fullof life atmosphere. This because I had the opportunity to met so many great people at theChemnitz University of Technology. I wish to thank:- Sabine, who become special friend of mine and for many discussions that we had, which arealways really enjoyable moments,- Steffen for friendships, constant support in any problems that I had and for the fun duringsailing,- Claudia, Mahendra, Cornell, Robert, Igor, Amit, Petra, Ina, Susi, Zhongjie, Ailin, Yana forthe fan in laboratory, during conferences and in any time we spend together, your friendshipsgladness my life,- Ms.Goldmann, Ms. Klaus, Dr. Baumann, Mandy, glass-blowers for the kind assistancewhich can not be unmentioned.It was great pleasure for me to spend last years in the Physical Chemistry Departmenttherefore I would like to thank former and present group members for the familiaratmosphere.The cooperation and financial support of BASF SE is gratefully acknowledged.I wish to thank all those directly and indirectly helped me during my thesis.I owe my parents much of what I have become. They have always been there for me as ansteady support. I thank them for their love, their support, theirs believe in me and theirassurance during the past twenty-nine years.I also would like to thank my brother Jarek, who always trust in me and support me in anythink that I have done.Last but not the least, I would like to express special thanks to my husband Dawid – for yourendless patience, endless love and encouragement when it was most required - I thank you somuch!

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Curriculum VitaeJoanna M cfel-Marczewski

born on 06. August, 1980in Bydgoszcz, PolandPolish citizen

08/2008 - Research Chemist,BASF Construction Chemicals, Trostberg, Germany

Education

06/2005 – 08/2008 PhD. Candidate (Physical Chemistry)Chemnitz University of Technology, Germany“Self Incompatible Solvent”

10/2004 – 06/2005 Scientific work at the Chemnitz University of Technology,Physical Chemistry, Germany,“Biomineralization”

10/2003 - 10/2004 Diploma work at the Technical University Hamburg-Harburg,Polymer Composites, Germany,“The interfacial interactions between Carbon Nanotubes andPolybuthylene terephtalate (PBT) in synthesized Nanocomposites”

10/1999 - 10/2004 Master - Engineer in Chemical Technology,University of Technology and Life Science in Bydgoszcz,previously Academy of Technology and Agriculture in Bydgoszcz,Poland,

09/1995 - 06/1999 High School Boleslaw Krzywousty High School, Nak o nad Noteci , Poland

Scholarships

04/2002 - 07/2002 Socrates Scholarship at the University of Siegen

Publications

„Effect of silicon fertilizers on silicon accumulation in wheat“Mecfel, J.; Hinke, S.; Goedel, W. A.; Marx, G.; Fehlhaber, R.; Baucker, E.; Wienhaus O.J. Plant Nutr. Soil Sci. 2007, 170, 769-772

„Preparation and characterization of PBT/carbon nanotubes polymer”Kwiatkowska, M.; Broza G.; M cfel, J.; Sterzy ski, T.; Ros aniec Z.Composites 2005, 5, 2

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Patents

„Self Incompatible Solvents“Mecfel, J.; Koltzenburg, S.; Goedel, W. A.EP 01721690.7

Posters:

Symposium „Molekulare Systeme“,Chemnitz University of Technology, Germany28th November 2007“Self Incompatible Solvents - A New Approach to Increase Solubilization Power” Mecfel, J.; Koltzenburg, S.; Goedel, W. A.

Makromolekulares Kolloquium, Freiburg, Germany, 29th February, 2008“Self Incompatible Solvents”Mecfel, J.; Koltzenburg, S.; Goedel, W. A.

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SelbstständigkeitserklärungHiermit erkläre ich an Eides statt, die vorliegende Arbeit selbständig und ohne unerlaubteHilfsmittel durchgeführt zu haben.

Joanna M cfel-Marczewski

Chemnitz, den

Documents

Self Incompatible Solvent - monarch.qucosa.demonarch.qucosa.de/fileadmin/data/qucosa/documents/6042/data/diss.pdf · Self Incompatible Solvent von der Fakultät für Naturwissenschaften