Boron-based Lewis Acids and Anions - An Experimental and Theoretical
Study
INAUGURALDISSERTATION
zur Erlangung des Doktorgrades
der Fakultät für Chemie und Pharmazie
der Albert-Ludwigs-Universität Freiburg im Breisgau
vorgelegt von Hannes Böhrer
aus Schweinfurt
2014
Die vorliegende Arbeit wurde im Zeitraum von September 2010 bis Mai 2014 am
Institut für Anorganische und Analytische Chemie an der Albert-Ludwigs-
Universität Freiburg im Breisgau unter der Anleitung von Prof. Dr. Ingo Krossing
angefertig.
Vorsitzender des Promotionsausschusses Prof. Dr. T. Koslowski
Referent Prof. Dr. I. Krossing
Korreferent Prof. Dr. H. Wegner
Datum der mündlichen Prüfung 30.06.2014
Ein Großteil von Kapitel 4 wurde bereits veröffentlicht (a) und Teile von Kapitel 3
zur Veröffentlichung eingereicht (b):
a) H. Böhrer, N. Trapp, H. Scherer, I. Krossing, Pure and Applied Chemistry
2012, 84, 2317.
b) H. Böhrer, N. Trapp, D. Himmel, I. Krossing, submitted 2013 to Dalton
Transactions.
Kapitel 4 basiert auf der oben genannten Publikation. Aus Gründen der
Einheitlichkeit wurden die Formatierungen des Originaltextes, die Nummerierung
der Verbindungen, Abbildungen, Tabellen, Schemata, Referenzen etc.
angepasst. Zusätzlich wurden noch Abbildungen, Diskussionen und neue
Ergebnisse eingefügt. Die Genehmigung zur Reproduktion der Publikation wurde
bei dem entsprechenden Verlag eingeholt. Auf Ergebnisse Dritter wurde zu
Beginn der jeweiligen Kapitel hingewiesen. Ergebnisse zu Arbeiten, welche
während meiner Diplomarbeit an der Albert-Ludwigs-Universität Freiburg (2010)
begonnen wurden und im Zuge der Promotion weitergeführt und vertieft wurden,
sind ebenfalls Teil dieser Dissertation.
Danksagung
Zuerst möchte ich meinem Doktorvater Prof. Dr. Ingo Krossing für die
interessante Aufgabenstellung danken, ebenso für die Freiheit diese
eigenständig zu interpretieren. Seine Begeisterungsfähigkeit und Motivation, die
er wie kein Zweiter weiterzugeben vermag, haben zum Gelingen dieser Arbeit
entscheidend beigetragen.
Ein generelles Dankeschön geht an alle, die mich vor und während der Zeit
meiner Promotion begleitet haben:
Prof. Dr. Hermann Wegner für die Übernahme der Korreferats und meiner
Drittprüferin Prof. Dr. Caroline Röhr.
Meinen Bachelorstudenten Anne Asmacher und Philippe Weis, deren
experimentellen Arbeiten teilweise in diese Doktorarbeit aufgenommen wurden.
Meiner Laborkollegin Julia Schaefer sowie ihrem Nachfolger Mario Sander für die
wunderschöne gemeinsame Zeit im Labor 330.
Dr. Nils Trapp für die Einführung in den Arbeitskreis und seine Hilfbereitschaft.
Den ehemaligen und aktuellen Mitgliedern des erweiterten Arbeitskreises
Krossing:
Anna Erle, Alexander Higelin, Alexander Rupp, Benedikt Burgenmeister,
Boumahdi Benkmil, Brigitte Jörger, Carola Sturm, Carsten Jenne, Christoph Bolli,
Christoph Schulz, Daniel Himmel, Daniel Kratzert, Dominic Kaase, Elias Frei,
Fadime Bitgül, Felix Brosi, Florian Stahl, Franziska Scholz, Gunther Steinfeld,
Harald Scherer, Heike Haller, Jana MacLaren, Jennifer Beck, Julia Klingele, Julia
Schaefer, Katharina Pütz, Kathrin Wagner, Lucia Alvarez Hernandez, Mara
Bürchner, Maribel Sierra Trillo, Marina Artamonova, Mario Sander, Marin
Lichtenthaler, Martin Lieder, Mathias Hill, Mathias Kamp, Mathias Keßler,
Meiping Liu, Melanie Sanchez, Michael Hog, Michael Rohde, Michel Panzer, Mira
Schulz, Miriam Schwab, Nils Trapp, Olaf Petersen, Pengcheng Zhang, Petra
Klose, Philipp Eiden, Przemek Malinowski, Rainer Riebau, Robin Brückner, Safak
Bulut, Sascha Goll, Sebastian Hasenstab-Riedel, Stefanie Reininger, Thomas
Vent-Schmidt, Timo Huxel, Tobias Engesser, Tobias Köchner, Tobias Schlöder,
Ulf Sachs, Ulrich Preiss, Valentin Dybbert, Valentin Radtke, Vera Brucksch,
Werner Deck, Witali Beichel,
Dr. Harald Scherer für die Einführung in die NMR-spektroskopische Messtechnik
und fortwährende Weiterbildung, sowie die unzähligen Stunden der Diskussion.
Dr. Daniel Himmel für die Unterstützung bei quantenchemischen Rechnungen.
Dr. Valentin Radtke und Dr. Philipp Eiden für die Hilfe bei der Aufnahme und
Analyse von cyclovoltammetrischen Messungen.
Dr. Nils Trapp, Dr. Daniel Kratzert, Boumahdi Benkmil, Olaf Petersen, Franziska
Scholz und Mario Sander für die Aufnahme und das Lösen der Kristallstrukturen.
Den Glasbläsern Tim Lecke und Martin Walter für das Anfertigen von
Glasgeräten sowie etwaiger Reparaturen.
Winni „Vollgas“ Weber für die Versorgung mit Gasen und die lustigen Gespräche.
Den Mädels aus dem Sekretariat, Brigitte Jörger und Vera Brucksch, für die
tatkräftige Unterstützung abseits des Labors. Ebenso Anita Weidner und Stefanie
Kuhl
Den Männer und Frauen aus der Werkstatt für die Reparaturen und die
Instandhaltung jeglicher Geräte.
Dem Team der Chemikalienausgabe für die Versorgung und Entsorgung von
Chemikalien.
Anne, Alex, Elias, Felix, Franzi, Heike, Julia, Mara, Mario, Matze (DJ) und Max
für die lustige Zeit abseits der Uni.
Der Mannschaft des SV Blau-Weiss Wiehre für den Ausgleich und die
gemeinsamen Erfolge.
Mein größter Dank geht an meine Familie und Stephanie für die Unterstützung in
jeglicher Hinsicht.
Table of Contents
List of Abbreviations IV
1 Introduction 1
1.1 Lewis Acids 1
1.2 Bifunctional Lewis Acids 6
1.3 Experimental and calculated Lewis acidity scales 10
1.4 Frustrated Lewis Pair (FLP)-Chemistry 13
1.5 Weakly Coordinating Anions (WCAs) 15
1.6 References 18
2 Objectives 25
2.1 References 26
3 The Lewis acid B(Ohfip)3 and innovative Lewis acid scales 27
3.1 Introduction 27
3.2 Results and Discussion 30
3.2.1 Synthesis and Characterization 31
3.2.2 Quantum Chemical Investigations 37
3.2.3 Analysis 46
3.3 Conclusion 48
3.4 Experimental Section 48
3.4.1 Theoretical Methods 48
3.4.2 Synthesis and Characterization 49
3.5 References 54
3.6 Appendix 27
4 Synthesis and Oxidation of Hexafluoroisopropoxyborylferrocenes 75
4.1 Introduction 75
4.2 Results and Discussion 78
4.2.1 Synthesis and NMR Spectroscopic Characterization 78
I
4.2.2 X-Ray Crystal Structures 80
4.2.3 Cyclo Voltammetric Characterization 85
4.2.4 Quantum Chemical Investigations 90
4.3 Conclusion 92
4.4 Experimental Section 93
4.4.1 Theoretical Methods 93
4.4.2 Synthesis and Characterization 93
4.5 References 97
4.6 Appendix 75
5 Approaching New Borate Based Weakly Coordinating Anions (WCAs) 103
5.1 Introduction 103
5.2 Results and Discussion 105
5.2.1 Quantum Chemical Calculations 105
5.2.2 Synthesis and Characterization 108
5.2.2.1 Attempts to synthesize [((CF3)3CO)2B(C6F5)2]− 8 110
5.2.2.2 Synthesis of [((CF3)2HCO)2B(C6F5)2]− 9 117
5.2.2.3 Synthesis of [(CF3CH2O)2B(C6F5)2]− 10 120
5.2.2.4 Metathesis and Alternative Synthesis Routes 124
5.3 Conclusion 129
5.4 Experimental Section 130
5.4.1 Theoretical Methods 130
5.4.2 Synthesis and Characterization 130
5.5 References 135
5.6 Appendix 137
6 Conclusion 141
6.1 References 145
7 Crystal Structure Data 146
II
8 Abstract 149
9 Kurzzusammenfassung 150
III
List of Abbreviations
a,b,c unit cell axes
Å Ångström
ant anthracen
ATR attenuated total reflection
b broad (NMR)
CCDC the Cambridge Crystallographic Data Centre
CIA chloride ion affinity
COSMO conductor-like screening model
COSY correlated spectroscopy
Cp cyclopentadienyl
CuD copper decomposition
d doublet (NMR)
d distance
D diffusion coefficient
DFB difluorobenzene
DFT density functional theory
E0 redox potential
E0’ formal potential
e.g. for example
Et ethyl
eq equivalent
IV
Fc ferrocene
FIA fluoride ion affinity
FLP frustrated Lewis pair
G Gibbs energy
H enthalpy
hfip CH(CF3)2
HIA hydride ion affinity
HMBC heteronuclear multiple bond correlation
HOMO highest occupied molecular orbital
HSQC heteronuclear single quantum correlation
ISS steady state current
J coupling constant
iPr iso-propyl
IR infrared
K Kelvin
k0 rate constant of the heterogeneous electron transfer
LA Lewis acid (Chapter 1)
LA ligand affinity
lat latin
LUMO lowest unoccupied molecular orbital
m Multiplett (NMR)
m medium (IR)
Me methyl
V
m meta
MIA methyl ion affinity
MO molecular orbital
NHC N-heterocyclic carbene
NMR nuclear magnetic resonance
NOE nuclear Overhauser effect
NOESY nuclear Overhauser effect spectroscopy
o ortho
p para
PD proton decomposition
Ph phenyl
ppm parts per million
g gaseous
q quartet (NMR)
quint. quintet (NMR)
r radius
reff effective radius
RF C(CF3)3
RT room temperature
S entropy
s strong (IR)
s singlet (NMR)
SCF self-consistent field
VI
solv solvation
t triplet
tBu tert-butyl
THF tetrahydrofuran
UME ultra micro electrode
UV/VIS ultraviolet/visible
vs very strong (IR)
vw very weak (IR)
w weak (IR)
WCA weakly coordinating anion
XRD X-ray diffraction
α* dip angle
α, β, γ unit cell angles
γ activity coefficient
δ chemical shift (NMR)
Δ difference
Δν1/2 half width
εr dielectric constant
θ scattering angle
μ absorption coefficient
VII
VIII
1 Introduction
An acid (lat.: acidum) was in alchemy a material prima (lat.: first material). The
term “acid” is derived from the probably oldest known acid, the vinegar (lat.:
acetum). The first acid-base concept was formulated by Arrhenius in the year
1887. It says, that in an aqueous solution, an acid dissociates in two particles, a
proton and an anion, and a base accordingly in a hydroxide ion and a cation.
However this definition is just valid in water and only considers acids and bases,
which separate H+ or [OH]−.[1]
J .N. Brønsted and T. Lowry 1923, from each other independently, defined an
additional acid-base concept, which is not constricted to the solvent water and
still used today.[2] In their concept they postulate, that an acid emits a proton to
water or another compound and as result a protonated base is generated (proton
donator). From this follows, that ions of an aqueous acid solution do not originate
from dissociation in an anion and a cation, but from a chemical reaction of an
acid with water. On the other hand the basic effects of a compound relies on
accepting protons of water molecules, which forms deprotonated acids, which in
turn are responsible for the basic character of a solution (proton acceptor).
AcidProton release
Proton acceptionBase + Proton
Equation 1.1: Brønsted’s acid-base definition.
1.1 Lewis Acids
Also in the year 1923, G. N. Lewis extended the acid-base theory by making it
independent of the proton:
“A basic substance is one which has a lone pair of electrons which may be used
to complete the stable group of another atom, and that an acid substance is one
which can employ a lone pair from another molecule in completing the stable
group of one of its own atoms.”[3]
1
Therefore, a Lewis acid is defined as a compound that acts as electron acceptor.
This includes compounds with incomplete electron octet (e.g.: BH3, BF3, AlCl3),
halogen compounds with unsaturated coordination (e.g.: PF5, SbF5, SiCl4) and
molecules with unsaturated double bonds (e.g.: SO2). Lewis acids form with
Lewis bases (electron donators) Lewis acid-base complexes. A Lewis base is
also a Brønsted base, but a Lewis acid does not need to be a Brønsted acid.
Lewis acidAssociation
DissociationLewis acid-base complexLewis base+
Equation 1.2: Lewis’ acid-base definition.
In those complexes electron density of a free electron pair in the HOMO of the
Lewis base is transferred into the energetically suitable leveled LUMO of the
Lewis acid, to reduce the electron deficiency of the Lewis acid via a polarized
dative bond. For his work to get a better understanding of chemicals bonds G. N.
Lewis was nominated 35 times for the noble prize, but he never got the honor to
receive the desired award.
Pearson’s HSAB concept: R. G. Pearson introduced in the year 1963 the
concept of hard and soft acids and bases (HSAB), to explain their affinity to each
another, which apparently not directly depends on the electronegativity or other
macroscopic properties. He postulated that hard Lewis acids prefer to combine
with hard Lewis bases and soft Lewis acid with soft Lewis acids.[4] Hard Lewis
acid are defined as small and hard to polarize electron acceptors, while soft
Lewis acids own a capacious acceptor atom and the valence electrons are
relatively easy to polarize.[5] Due to the high effective nuclear charge and the
resulting strong electrostatic attraction of the valence electrons, hard compounds
are extremely difficult to polarize and the molecular orbitals are positioned on a
lower energetic level. Scheme 1.1 shows that the combination hard/hard is
electrostatically dominated, whereas the soft/soft pair is frontier orbital controlled
and possesses a mainly covalent bond.
2
E
Acid Base Acid Base
σ
σ∗
σ
σ∗
hard/hard soft/soft
Scheme 1.1: MO charts for the dative bond between a hard acid and a hard base (left) as well as between a soft acid and a soft base (right).
Classical Lewis acids: Classical Lewis acids are e.g. trigonal planar compounds
with an empty pz orbital. Among others e.g. the boron trihalides are classical
Lewis acids. At room temperature BF3 and BCl3 are gaseous, BBr3 is a liquid and
BI3 is already solid. The halide atoms in these compounds are capable to donate
electron density from their filled nonbonding pz orbital to the empty pz orbital at
the boron atom, which increases the strength of the B-X bonds. The larger the
overlap of the pπ-orbitals the stronger is the back-donation. Additionally to the π-
donation, especially the B-F bond is enhanced by electrostatic interactions. That
makes the B-F bond in BF3 to one of the strongest bonds known.[6] Despite this
backbonding, BF3 forms readily adducts with a variety of Lewis bases and has
found a multitude of applications (see below), whereas adducts of BBr3 and BI3
are relatively unstable. Due to their weaker B-X bonds, halide displacement
reactions prevail. It was stated that the different degree of π-donation is the main
criterion for the Lewis acidity, since it will get lost on pyramidalization. At least the
increasing sequence of Lewis acidity BF3 < BCl3 < BBr3 < BI3 is in contrast to the
expected order according to electronegativity arguments. But for this back-
bonding explanation is little compelling evidence, since it has been shown, that
the overlap between the boron and chlorine pπ orbitals in BCl3 is even larger than
the overlap of the respective boron and fluorine orbitals in BF3.[7] A further
approach figures out, that the halogen ligands remain close-packed throughout
the adduct formation and that the bond lengths increase accordingly.
Consequently, lengthening the strong B-F bond takes more energy than
3
lengthening the weaker B-Cl bond.[8] Another approach is based on the frontier
orbital model of chemical reactivity.[9] The LUMO energy of BCl3 is lower than the
LUMO energy of BF3, and hence the interaction with an occupied donor orbital is
stronger and there is a larger covalent bond part. However, the former given
sequence of Lewis acidity only counts for Lewis bases, which are strong enough
to change the geometry to a tetrahedral one. In contrast, for weak Lewis bases
like CO, HCN, CH3CN and CH3F, the order is reverse and follows the one on the
basis of the charge on the boron atom expected.[10]
The trivalent Lewis acid tris(pentafluorophenly)borane is obtained, when the
halogen ligands in BX3 compounds are replaced by the pentafluorophenyl group,
which represents one of the most present Lewis acids nowadays.[11] Its Lewis
acidity is in the same range as that of BF3, considering experimental scales.[12, 13]
One big advantage of B(C6F5)3 compared to the borontrihalides is the increased
stability in the presence of protic agents, since the B-C and C-F bonds are more
inert.[11, 14] Although the bulkier residues in B(C6F5)3 hamper the access to the
Lewis acid centre and a little Lewis acidity is lost, the compound is more easily
handled, without the problems associated with the reactive B-X bonds.[15]
From all these approaches follows that determining the Lewis acidity of a
compound is not trivial. Orbital overlap, reorganization energy,[16] LUMO energy,
electronegativity and steric effects of a Lewis acid are involved, but even more,
the strength of a Lewis acid depends on the bonding partner, the Lewis base.
Defining and establishing an accurate Lewis acidity scale is part of this work and
elaborated in more detail later on.
Application of classical Lewis Acids: Borontrifluoride is the most widely used
boron halide.[17] Especially in organic synthesis it finds a variety of applications.
BF3 is used as catalyst in Friedel-Crafts alkylation reactions[18], to catalyze the
cleavage of ethers to alcohols[19] and in the nitration[20] of aromatic compounds
(e.g. as BF3•2CF3CH2OH complex).[21] Also BF3 is used in conjunction with a
proton donor as an initiator in olefin polymerization reactions. Every year
approximately 2300 – 4500 tons of BF3 and its adducts are produced.[17] BCl3
finds similar applications like BF3 as Lewis acid catalyst in the organic synthesis,
although not as much, since it is more prone to halide displacement reactions. A
4
main application of BCl3 is the preparation of boron fibers. Usually a tungsten
wire is passed through a reaction chamber filled with BCl3 and hydrogen. BCl3
gets reduced and elementary boron deposits on the wire, building a strong boron
fiber.[22] The more complex AlCl3 is likely the most commonly used Lewis acid.
Especially in Friedel-Crafts type reaction AlCl3 becomes indispensable.[23]
Furthermore it is capable to carry out degradation reactions of perfluorinated
compounds,[24] electrophilic aromatic substitutions,[25] and normal and inverted
electron-demand Diels-Alder reactions,[26] just to name a few of a bunch of
applications.[27] Tris(pentafluorophenly)borane finds applications in silylation of
alcohols[28], Si-H bond activation[29] or reductive cleavage of alcohols and
ethers.[30] In inorganic chemistry it is used for anion-binding[31] or for synthesis of
weakly coordinating anions.[32, 33] Since the introduction of frustrated Lewis pair
chemistry one of its most important use is as component in FLP systems (see
Chapter 1.4).[34, 35]
Non-classical Lewis acids: The hypercoordinated compound SbF5 is the
strongest conventional Lewis acid. It is able to generate reactive noble gas
cations like [XeF]+, [Xe2F3]+,[36] [XeF3]+[37] or [XeO2F]+ by fluoride abstraction with
SbF5 giving [SbF6]− or [Sb2F11]−.[38] Furthermore SbF5 was used to define Lewis
superacidity (see 1.3).[39] Carbon and silicon form carbo-cations or silylium-
cations that are isoelectronic to BX3 compounds, but much harder to design,
since they are highly reactive ions, which need to be stabilized by inert weakly
coordinating anions (WCA, see 1.5). A neutral molecule with a Lewis acid atom of
the group 14 elements is a curiosum in Lewis acid chemistry, but was realized in
the silylium zwitterion SiMe2CH2CB11Cl11.[40] The environment of the Si atom is
only modestly pyramidalized, which makes this compound a powerful Lewis acid.
A classification of the Lewis acidity of the silylium zwitterion and a variety of
common Lewis acids is given in chapter 3.
5
1.2 Bifunctional Lewis Acids
Multifunctional Lewis bases are often also named as multidentate Lewis bases,
since their free electron pair is compared with a tooth. However Lewis acids do
have rather an electron cavity, so this terminology is inept for Lewis acids and in
the following Lewis acids with two Lewis acid atoms are addressed as
bifunctional.
Multifunctional Lewis Acids: As mentioned in 1.1, Lewis acids are used in
organic synthesis to activate functional groups. The improvement of reaction
conditions can be enhanced by a second Lewis acid molecule, which coordinates
at the same Lewis base position. However adducts of two independent
monofunctional Lewis acids are entropically and kinetically disfavored. To enable
double coordinated adducts, both Lewis acid atoms are linked to a bifunctional
Lewis acid (Figure 1.1). The chelate effect reduces the entropy term and the
kinetic inhibition is reduced.[6] The abilities of a given Lewis acid are strongly
influenced by the type of bridge connecting the Lewis acid atoms. On the one
hand the electronic properties of the linker are important for the Lewis acidity, on
the other hand its geometry determines the bite angle of the resulting Lewis acid-
base complex and thus the type of Lewis bases to be bound.[41]
OOO
LA LALALA LA
Figure 1.1: Different coordinated carbonylfunctions (left: one monofunctional LA; mid: two monofunctional LA; right: a bifunctional LA).
Because of the high reactivity of multifunctional Lewis acids and, as a
consequence, the synthetic difficulties associated with it, the development lagged
far behind the chemistry of multifunctional Lewis bases, like e.g. crown ethers.[42]
The common strategy to stabilize multifunctional Lewis acids is the incorporation
of ligands, which are able to reduce the Lewis acidity through σ-donation or
π-bonding. However, this is in contrast to the actual aim of this class of
compounds and the degree of stabilization has to be considered.
6
Evolution of boron based bifunctional Lewis acids: Pioneering works on
bifunctional Lewis acids were done by Shriver and Biallas with the synthesis of
1,2-bis(difluoroboryl)ethane (Figure 1.2 a) and some adducts.[43] Despite this
early and promising result no further progress was made for the next twenty
years until the “hydride sponge” 1,8-bis(dimethylboryl)naphthalene (Figure 1.2 b)
was introduced.[44] This “sponge” is capable to bind H− and also other anions like
F− and OH−. It was shown by competitive binding studies, that bifunctional Lewis
acids bind these anions better than a monofunctional Lewis acid. Two years later
1,2-(BMe2)2C6H4 (Figure 1.2 c) was published, but it is thermally unstable and
condenses intermolecular under BMe3 elimination to 9,10-diboraanthracene
compounds.[45] To avoid this decomposition, 1,2-(B(C6F5)2)2C6F4 was synthesized
later on, additionally the Lewis acidity is enhanced by substitution of the H atoms
of the benzene ring with F atoms.[46] Further interesting attempts were done by
Wrackmeyer and Siebert,[47] by developing cis-1,2-diboryl compounds (Figure 1.2
d), which are capable of binding small ions (but not the trans isomers). The
Achilles’ heel of these compounds is the occurring UV-isomerization to a
equilibrium of cis and trans isomers, since the empty p orbital of the boron atom
is able to overlap with the C-C π-bond and destabilize the double bond.[48]
F2BBF2
BMe2BMe2
(C6F5)2BHN
B(C6F6)2
BEt2
EtEt
Et2B
BMe2
BMe2
Si(C6F5)2B(C6F5)2B
a cb
d fe
Figure 1.2: Selected examples of boron based bifunctional Lewis acids.
Compounds with a significantly reduced Lewis acidity are for instance
HN(B(C6F5)2)2 (Figure 1.2 e) with a B-N bond possessing a large π-part,[49] or the
1,3-diboryl Me2Si(CH2B(C6F5)2)2 (Figure 1.2 f) with a silicon atom in beta
position.[42, 50] Metallocenes, e.g. ferrocene could also be used to link two Lewis
7
acid atoms. The resulting ferrocenylboranes are subject of chapter 4 and
discussed there in more detail.
State of the art of bifunctional Lewis acids: The chemistry of mercury based
bifunctional Lewis acids is dominated by Wuest’s 1,2 diphenylenedimecury
compounds like 1,2 diphenylenebis(trifluoroacetato-O)dimercury (Figure 1.3 a).[51]
Their well defined orientation of the Lewis acid centers enables unique
coordination chemistry, like quadruple coordination of carbonyl groups by two
bifunctional Lewis acids.[52] Interestingly examples for aluminum based
bifunctional Lewis acids are the (2,7-disubstituted-1,8-
biphenylenedioxy)bis(dimethylalumium) compounds (Figure 1.3 b),[53, 54] which
are able to catalyze Claisen rearrangement reactions,[54] Michael additions,[54]
Meerwein-Pondorf-Verley reductions[55] or activate ether functionality.[54] Further
examples are 2,2’-(1,2 or 1,3 substituted)bisphenols[56] or derivates of
2,7-bis(1,1-dimethylethyl)fluorine-1,8-diol (Figure 1.3 c).[57]
a cb
Hg
Hg
OOCCF3
OOCCF3
O OAlR2R2Al O O AlEt2AlEt2
Figure 1.3: Mercury and aluminum based bifunctional Lewis acids.
One of the main research objectives in the bifunctional Lewis acid chemistry is
anion sensing, especially of the potentially toxic fluoride ion. Actually, the trend is
towards heteronuclear bifunctional Lewis acids. Figure 1.4 a shows a B/Hg
bifunctional compound, which readily and selectively chelates F− anions in a
variety of solvents and in the presence of Cl−, Br−, I− or CN− anions.[58] However,
the trick of this compound is that it acts as a phosphorescent anion sensor, since
the population of the empty pz orbital by the lone pair of fluoride leads to the loss
of conjugation and the naphthalenediyl chromophore phosphoresces. Another
example for heteronuclear fluoride anion binding are
o-(fluorosilyl)(dimesitylboryl)benzenes (Figure 1.4 b), which represent B/Si
bifunctional Lewis acids.[59] The chelate effect in these types of compounds can
be combined with favorable Coulomb interactions to strengthen the Lewis acidity
and the host-guest interaction, like in a B/P zwitterionic Lewis acid
8
(Figure 1.4 c)[60] or a naphthalene based B/Hg Lewis acid (Figure 1.4 d)[61]. Both
compounds are still highly fluoride selective over other anions. Later on Gabbaϊ
published a series of heteronuclear cationic Lewis acids based on a bifunctional
phosphonium/borane Lewis acid (Figure 1.4 e).[62] In the advanced
[o-(Ph2MeSb)(Mes2B)C6H4]+ the Lewis acidity was further enhanced, since the
stibonium ion is larger, more electropositive and heavier than the phosphonium
ion.[63] From the same effects the analogous telluronium ion benefits.[64] To proof
the concept this “onium based strategy” was extended and the sulfonium
fluorosilane Lewis acid [1-Ant2FSi-2-Me2S-(C6H4)]+ (Figure 1.4 f) was obtained.[65]
This compound is still viable to complex fluoride but its Lewis acidity is not as
strong as in heteronuclear cationic borane based bifunctional Lewis acids.
B HgMes
MesC6F5
FSi
BMes Mes
F
R
R
F R = Me, PhB
P
MesF
Me Ph
B HgMes
Mes
F
NMe3
B PMes
Mes
Ph
PhMeF
SiS
F
Ant Ant
Me
MeF
a cb
d fe
Figure 1.4: Difunctional heteronuclear neutral and cationic Lewis acids and their fluoride chelate complexes (blue).
In the year 2010 the Wegner group presented the first inverse electron demand
Diels-Alder reaction (IEEDA) of 1,2-diazenes by a bifunctional Lewis acid.[66]
Long time the high LUMO level of 1,2-diazines hindered an extended IEEDA
chemistry of these compounds, since harsh reactions conditions were necessary.
By using a bifunctional Lewis acid as catalysts, the reaction can be performed
under milder conditions at lower temperatures. The general concept is based on
the simultaneous coordination of a bifunctional Lewis acid to both nitrogen atoms
of the 1,2-diazine. This leads to a lower LUMO energy and a reduced electron
density in the previously aromatic system, which ease the [4+2] cycloaddition
9
step. By the following release of N2 the aromaticity is recovered (Scheme 1.2),
the desired aromatic product is obtain and the bifunctional Lewis acid is
regenerated. This reaction type can be used for a variety of enols and enamines
to obtain 2,3-substituted napthalenes, which are hard to prepare by other
ways.[66-68] Furthermore cyclopropanated napthalenes can be obtained in a
diastereoselective kind by a followed sigmatropic rearrangement.[69]
BR2 BR2
NNNN
BR2 BR2
R2R3
NN
BR2 BR2
R3 R2
R3R2
N2
R1
R1
R1
R1
Scheme 1.2: Proposed catalytic cycle for the IEEDA reaction of 1,2-diazines catalyzed by a bifunctional Lewis acid.[67]
1.3 Experimental and calculated Lewis acidity scales
Naturally tabulating the strengths of Lewis acids is very important and useful to
estimate the potency of a given Lewis acid. But in contrast to the strength of
Brønsted acids, which are typically measured experimentally and ranked within
one homogenous medium on the basis of the well-known pH and pKa scales that
10
can be set absolute by using the correct reference state and anchor points, the
strength of Lewis acids depends on the formation of a Lewis acid-base pair.[70] In
this respect Brønsted acidity is a special case of Lewis acidity, in which only one
type of Lewis acid (the proton) interacts with a great variety of Lewis bases free
of choice. Thus, it is only possible to determine the absolute strength of a Lewis
acid, with respect to a well-defined Lewis base.
Experimental approaches to estimate the Lewis acidity: 1961 Lappert et al.
presented one of the first attempts to measure the relative acceptor strength of
Lewis acids.[71] Based on the coordination of ethyl acetate (as reference) to some
main group three and four halides, iron (III) chloride and a few organoboron
chlorides, they measured the vibrational stretching frequencies of the adduct to
obtain information about the bond lengths. The stronger the bond the higher is
the stretching frequency.[72] Approaches based on infrared carbonyl stretching
frequencies were continued by other groups with partly different Lewis bases,
e.g. measuring of the C-N bond strength in complexes as reference.[13]
Another concept to establish a quantitative scale of Lewis acidity based on
experimental data were introduced by Drago et al. with the double scale enthalpy
equation:[73]
−ΔH = EAEB + CACB
Equation 1.3: Double scale enthalpy equation.
ΔH is the enthalpy of a Lewis acid/base reaction measured in the gas phase or
poorly solvating media like alkanes or CCl4. The Lewis affinity of a Lewis acid A is
determined by two empirically measured parameters, EA and CA (E stands for
electrostatic and C for covalent) and respective by EB and CB for a Lewis base. A
disadvantage of this method is the need of a huge thermodynamic data set,
which is required to ascertain the correlation.
A more generally concept to rate the strength of Lewis acids was presented by
Larson and McMahon.[74] They determined the binding energies in the gas phase
of fluoride and chloride ions to a variety of Lewis acids with ion cyclotron
resonance halide-exchange equilibrium techniques and used these
thermochemical data to construct scales of Lewis acidity. Additionally other
11
groups showed, that the fluoride affinity can be determined experimentally by
calorimetry or effusion techniques.[75]
Furthermore, there are two decent scales of relative Lewis acid strength of
borane Lewis acids based on NMR spectroscopy. The first one, the Gutmann-
Beckett method, makes use of the change of the 31P NMR shift of
triethylphosphine oxide upon complexation. Gutmann introduced the acceptor
number (AN) scale of solvents, by setting two anchor points relating to the 31P
NMR shift of Et3PO in the weak Lewis acid solvent hexane (AN = 0) and in the
very strong Lewis acid solvent SbF5 (AN = 100).[76] 1996 Beckett et al. extended
this former scale for solvents to general Lewis acids.[77] The second order of
Lewis acids was introduced by Childs et al. and is based on the chemical shift
difference of the H3 proton in crotonaldehyde upon complexation of a Lewis acid,
since the magnitude of the shift of this proton is the highest. [78]
P
OLA
OLA
H3
Figure 1.5: Complexes used to determine the strengths of Lewis acids by the change of chemical shifts in NMR spectroscopy.
Disadvantages of experimental methods to measure Lewis acidity are the large
uncertainty and the dependency on solvents. Furthermore, it is very hard to
determine highly reactive molecules, since they cannot be stabilized or may
decompose during measurement.
Quantum chemical approaches to estimate the Lewis acidity: As mentioned
before, it is only possible to determine the absolute strength of a Lewis acid, with
respect to a well-defined Lewis base. Towards this aim, in 1984 the fluoride ion
affinity (FIA) was introduced by Bartlett et al. to classify the strength of the Lewis
acid A by the enthalpy that is released by binding a fluoride ion.[79] This concept
was continued by many others.[16, 80-83] To avoid the problems that appear by the
calculation of a “naked fluoride ion”, this approach was improved by using the
experimental FIA of OCF2 of 209 kJ mol−1 as an anchor point in a
(pseudo-)isodesmic reaction.[84]
12
COF3− + A AF− + COF2
OCF3−COF2 + F−
A + F AF−
DFT-Method
experimental
−ΔH = FIA
Equation 1.4: Christie’s definition of the fluoride ion affinity (FIA).
Using this method, the strengths of many Lewis acids were classified.[39]
However, the fluoride ion is a hard base in the HSAB sense (see chapter 1.1), so
the FIA may be deceptive for HSAB soft Lewis acids. Thus, it is advisable to
compare methods with different approaches to find a convenient Lewis acid.
Therefore, analogously to the FIA, the affinities to ions of varying hardness (i.e.,
Cl−, H− and CH3− = Me−) were used to rank the strengths of Lewis acids.[83, 85]
Unfortunately, there is no consistent reference system, which would allow
comparing trends between the different methods and lead to a broadly applicable
Lewis acid scale. FIA values were also used to define Lewis superacidity. Every
Lewis acid with a higher FIA than SbF5 (489 kJ mol−1) is classified as Lewis super
acid.[39]
1.4 Frustrated Lewis Pair (FLP)-Chemistry
In the year 1942 the group of Brown discovered a system, which does not follow
the classical Lewis acid-base chemistry.[86] Besides other pyridines, they
investigated the reaction behaviors of 2,6-lutidine with BF3 and BMe3, but
observed only a adduct formation for treatment with BF3, but not for BMe3
(Scheme 1.3).
N BF3BF3 BMe3
N NMe3B
Scheme 1.3: Unexpected behavior of 1,2-lutidine with bulky Lewis acids.
13
By using molecular models (Figure 1.6) they attributed the failed adduct formation
to steric interactions of the o-methyl groups. They continued studying this
phenomenon, but did not test the 2,6-lutidine/BMe3 mixture on reactivity.[34, 87]
Figure 1.6: Molecular models of the coordination compounds of 2,6-lutidine with BF3 (left) and BMe3 (right).[86] Reprinted (adapted) with permission from H. C. Brown, H. I. Schlesinger, S. Z. Cardon, Journal of the American Chemical Society 1942, 64, 325. Copyright 1942 American Chemical Society.
In 1959 it was reported, that 1,2-didehydrobenzene with triphenylborane and
triphenylphosphine forms an o-phenylene-bridged phosphonium-borate.[88] Later
on the anionic polymerization of buta-1,3-diene with the trityl anion and BPh3 did
not give the expected polybutadiene compound, but the trapping product as ate-
complex.[89] Stephan’s group took up those results and laid the foundation for
chemistry associated with frustrated Lewis pairs (FLP) by observing the
reversible activation of hydrogen by a frustrated Lewis pair (Scheme 1.4).[34, 90]
The air and moisture stable phosphonium borate salt releases hydrogen upon
heating above 100°C.
F F
FF
(Me3C6H2)2P B(C6F5)2H2 25°C
-H2 150°C
F F
FF
(Me3C6H2)2P B( 6F5)2C
H
H
Scheme 1.4: Reversible H2 activation.
Hydrogen activation: Apart from intramolecular FLP chemistry (Scheme 1.4),
three-component FLP chemistry is moved into focus, since the facile formation of
phosphonium borates from the reaction of sterically demanding phosphanes,
boranes and hydrogen and the subsequently heterolytic cleavage of hydrogen is
14
more simple. It is possible to cleave hydrogen by tert-butyl-phosphine and
B(C6F5)3 at room temperature, where steric congestion precludes the
neutralization reaction via adduct formation.[91] In contrast to the zwitterionic
(C6H2Me3)2PH(C6F4)BH(C6F5)2 no liberation of H2 occurs at heating to 150°C.
B(C6F5)3 + tBu3PH2
1 atm, 25°C[tBu3PH][HB(C6F5)3]
Scheme 1.5: Three-component heterolytic cleavage of hydrogen by tert-butyl-phosphine and B(C6F5)3.[91]
Further examples for the widely spread used phosphorus/boron FLP system to
activate hydrogen are alkenylene-linked[92] or bisposphino naphthalene based
FLP systems.[93] Furthermore, pairs of N-heterocyclic carbenes (NHCs)[94] and
boron (C/B) or amines[95] or pyridines[96] with boron pairs (N/B) have been
investigated.
Applications of FLP systems: Today’s main applications of FLP chemistry are
metal-free hydrogenation catalyst reactions. First attempts in hydrogenation were
done simultaneous with the investigation of the first FLP systems, which is for
sure owed to their analogy to Noyori-hydrogenation catalysts.[90, 97] Since this
topic of FLP chemistry is only a few years old, the progress is impressive. Imines,
aziridines, enamines, silyl enol ethers, diimines, metallocene derivates and
nitrogen based heterocycles can be reduced by FLP systems.[98] An extensive
overview of FLP catalyzed hydrogenations is given in recent reviews.[98, 99]
Further applications are ethene/alkyne exchange reaction,[100]
1,2-addition reactions to carbonyl compounds[101] and activation of small
molecules like alkenes[102] or CO2.[103]
1.5 Weakly Coordinating Anions (WCAs)
Highly electrophilic cations are usually only observed in the gas phase in a mass
spectrometer or an argon matrix near absolute zero, since in condensed phase
they tend to coordinate or decompose. To investigate such reactive cations under
laboratory conditions, the abilities of the counter-ion are decisive. It has to be
extremely stable and the interactions between cation and anion should be very
weak, otherwise, the geometry of the cation is distorted. So the class of weakly
15
coordinating anions (WCAs) was invented, which are capable to stabilize reactive
cations under “pseudo gas phase conditions”.[32, 104]
Properties of WCAs: To prevent decomposition, WCAs benefit from a non-
nucleophilic and chemically robust surface area (e.g C-F bonds). The charge of a
WCA should be monovalent and delocalized equally over the whole molecular
ion. So the larger the ion, the higher is the potential degree of delocalization.
However. the most basic site represents the weak point of WCAs, since it is the
primary location of electrophilic attack. Mostly this spot is an electronegative atom
and can be sterically shielded by bulky ligands. The large volume of the anion is
accompanied by minimization of Coulomb interactions, which leads to lower
lattice energies and leads to high solubilities.[32, 104]
Applications of WCAs: Weakly coordinating anions (WCAs) have become an
indispensable tool in chemistry and found versatile applications. For example in
reactive cations e.g. [CX3]+ (X = Cl, Br, I),[81, 105] [TeX3]+ (X = Cl, Br, I),[106]
[N5]+,[107] [P9]+,[108] [AuXe4]2+,[109] [H(Et2O)2]+,[110] [Zn2]2+,[111] benzidine radical
cations,[112] H+, CH3+, Me3Si+,[113] [Cu(S12)(S8)]+,[114] the two-coordinate gallium ion
[tBu3Si-Ga-SitBu3]+,[115] triarylsilylium or -germylium ions,[116] protonated
benzene,[117] the 2-norbornyl cation,[118] the ion-like Me3Si-F-Al(OC(CF3)3)3 Lewis
acid,[119] ethane-copper,[120] -silver[121] or -gold complexes.[122] Beside the
preparation of reactive cations WCAs found applications in boron[123] or
aluminum[124] based lithium ion batteries, catalytic C-F activation,[125] ionic
liquids[126] and electrochemistry.[127]
Relative stability of WCAs: First attempts to order the relative stability and the
coordinating abilities of various WCAs were performed by Reed et al., using the 29Si NMR shift of the Si(iPr)3X (X = WCA) silylium ion pair as a measure, whereby
shifts to lower field indicate a more pronounced cationic character of the silylium
cation which is an evidence for a less coordinating anion X−.[128] Later, the same
group reported the N-H stretching frequencies of tri-n-octylammonium salts
(n-oct)3NH+X– (X = WCA) in CCl4 for a series of weakly basic anions. The
fundamental idea is that the weaker the NH+-X− interaction, the stronger is the
N-H bond and the higher the frequency of the NH stretching vibration.[129]
16
In independent work, the Krossing group judged the relative stabilities of WCAs
based on calculations on a representative set of WCAs and their parent Lewis
acids.[82] This method allowed to order the relative stabilities and coordinating
abilities of WCAs by the Ligand affinity (LA), decomposition in the presence of a
hard (proton decomposition PD) and a soft electrophile (copper decomposition
CuD), the position of the HOMO, the HOMO-LUMO gap and the before
mentioned FIA, which is capable to predict the stability of a given (WCA) by
viewing the FIA of its parent Lewis acid. The higher the FIA of the Lewis acid, the
more stable is the respective WCA towards ligand abstraction.
M(L)n− + L−
HL+ H+
+ Cu+ CuL−ΔH = CuD
M(L)n−
M(L)n−
ΔH = LA
−ΔH = PD
M(L)n−1
M(L)n−1
M(L)n−1
+
+
Equation 1.5: Definition of Ligand affinity (LA), proton decomposition (PD) and copper decomposition (CuD).
Since a single methodology was chosen for calculations the relative ordering of
the stabilities of all WCAs and their underlying Lewis acids are, due to error in
calculation most likely correct.
17
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Organometallics 2012, 31, 7298. [116] A. Schäfer, M. Reißmann, S. Jung, A. Schäfer, W. Saak, E. Brendler, T.
Müller, Organometallics 2013, 32, 4713. [117] F. Scholz, D. Himmel, L. Eisele, W. Unkrig, I. Krossing, Angewandte
Chemie-International Edition 2014, 53, 1689. [118] F. Scholz, D. Himmel, F. W. Heinemann, P. V. Schleyer, K. Meyer, I.
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2 Objectives
Lewis acids are indispensable in chemistry. Due to their widespread applications,
ranking Lewis acidity is a powerful tool to rate the abilities of a given Lewis acid.
However, the strength of a Lewis acid cannot be determined like the strength of a
Brønsted acid, since Brønsted acidity is a special case of Lewis acidity, in which
only one type of Lewis acid (the proton) interacts with a great variety of Lewis
bases free of choice. The strength of a Lewis acid can only be evaluated with
respect to a well-defined Lewis base. Many promising quantum chemical
approaches have been made to obtain Lewis acidity values, using F−, Cl−, H− or
methyl anions as different Lewis bases.[1] Unfortunately, there is no consistent
reference system, which would allow comparing trends between the different
methods and lead to a broadly applicable Lewis acid scale.
One objective of this thesis was to develop and validate a unified Lewis acid
scale, so that the relative values of the Lewis acidity towards different bases are
better comparable. Using this method the most common Lewis acids should be
ranked with respect to their FIA (Fluoride Ion Affinity) CIA (Chloride Ion Affinity),
HIA (Hydride Ion Affinity) and MIA (Methyl Ion Affinity). In this context the known
scales for the WCA stability can be expanded with the PD and CuD values of a
series of hitherto not explored WCAs, which are based on the investigated Lewis
acids.[2]
A further objective was to test the suitability of the Lewis acid B(Ohfip)3 (Ohfip =
OCH(CF3)2) combined with different sterically encumbered phosphanes, amines,
pyridines and NHCs (N-Heterocyclic carbenes) to form FLP systems capable to
cleave hydrogen. In this connection the crystal structures of B(Ohfip)3 and
possible adducts should be revealed.
Ferrocenylboranes are a very attractive class of compounds for a number of
reasons. Their special bonding situation, electrochemical properties and the
Lewis acidity made them subject of many investigations.[3] Crystal structures of
their oxidized ferrocinium derivates are still rare.
An emphasis of the present thesis was put on synthesis and extended
characterization (X-ray diffraction, NMR, IR, RAMAN, UV/VIS and CV), of
25
26
[1-(BOhfip)2Fc] and [1,1’-(BOhfip2)2)Fc], as well on their oxidized species.
Especially exploring the crystal structures is of importance to deepen the
knowledge of ferrocenylboranes and the electronic and steric behavior of the
Ohfip residues in boron-based Lewis acids. Quantum chemical investigations
complete the analysis.
Since WCAs have found application in a variety of different ways in chemistry,
the design and synthesis of new WCAs is more important than ever. Herein the
synthesis of heteroleptic WCAs, based on different sterically demanding alcohols
to give [(RFO)2B(C6F5)]− (ORF = OC(CF3)3, OCH(CF3)2, OCH2CF3), should be
investigated and improved. To gain an overview over the stability, quantum
chemical calculations were performed.
2.1 References
[1] T. E. Mallouk, G. L. Rosenthal, G. Muller, R. Brusasco, N. Bartlett, Inorganic Chemistry 1984, 23, 3167; J. W. Larson, T. B. McMahon, Journal of the American Chemical Society 1984, 106, 517; K. O. Christe, D. A. Dixon, D. McLemore, W. W. Wilson, J. A. Sheehy, J. A. Boatz, Journal of Fluorine Chemistry 2000, 101, 151; F. Bessac, G. Frenking, Inorganic Chemistry 2003, 42, 7990; J. C. Poutsma, O. E. Schroeder, R. R. Squires, Chemical Physics Letters 2004, 389, 433; L. O. Müller, D. Himmel, J. Stauffer, G. Steinfeld, J. Slattery, G. Santiso-Quinones, V. Brecht, I. Krossing, Angewandte Chemie-International Edition 2008, 47, 7659; A. Y. Timoshkin, G. Frenking, Organometallics 2008, 27, 371; D. J. Grant, D. A. Dixon, D. Camaioni, R. G. Potter, K. O. Christe, Inorganic Chemistry 2009, 48, 8811; L. A. Mück, A. Y. Timoshkin, G. Frenking, Inorganic Chemistry 2012, 51, 640.
[2] I. Krossing, I. Raabe, Chemistry-a European Journal 2004, 10, 5017. [3] M. Scheibitz, M. Bolte, J. W. Bats, H. W. Lerner, I. Nowik, R. H. Herber, A.
Krapp, M. Lein, M. C. Holthausen, M. Wagner, Chemistry-a European Journal 2005, 11, 584.
3 The Lewis acid B(Ohfip)3 and innovative Lewis acid scales
Philippe Weis performed experiments to FLP chemistry as part of his bachelor
thesis under my supervision. Peter Reiser performed some orientating
calculations under the supervision of Dr. Nils Trapp. These preliminary results
were reworked and led to about 20% of the Lewis affinity values shown in Table
3.5. The technical data for all executed calculations is deposited on the file server
of the group of Prof. Krossing (path: public\Ehemalige\Hannes Böhrer
2014\SI_Disseratation\SI_Tech_Chapt3.pdf).
3.1 Introduction
In the past years Lewis acids, and especially their strongest representatives,
were of great interest and found applications in catalysis, ionization,
rearrangement reactions, cycloaddition reactions and bond heterolysis
reactions.[1] Currently frustrated Lewis pair (FLP) chemistry begins to shine, since
FLPs are metal-free and capable to activate small molecules like H2, alkenes,
aldehydes and CO2.[2] Thus ranking Lewis acidity is very useful and a powerfully
tool to evaluate the abilities of a given Lewis acid. Unfortunately, the strength of a
Lewis acid cannot be determined like the strength of a Brønsted acid, since
Brønsted acidity is a special case of Lewis acidity, in which only one type of
Lewis acid (the proton) interacts with a great variety of Lewis bases free of
choice. That means that Brønsted acidity can be experimentally measured and
ranked in respect to the proton via the well-known pH and pKa scales. By using
the correct reference state and anchor points (the chemical potential of the
proton) the Brønsted acidity can be calculated and compared in different media.[3]
The experimental and as well the calculated data for Brønsted acidity can be set
absolute. But in contrast, the strengths of Lewis acids depend on the formation of
a Lewis acid-base pair and thus, it is only possible to determine the absolute
strength of a Lewis acid, with respect to a well-defined Lewis base. First
promising attempts to establish a method to measure the strength of a Lewis acid
were done in 1984 by the Bartlett group that used the fluoride ion as Lewis
base.[4] In this approach, the enthalpy that is released upon binding a fluoride ion
27
is used and connected with the strength of the Lewis acidity. The more negative
the enthalpy, the higher the fluoride ion affinity and the stronger the Lewis acid.
The concept of fluoride ion affinity was continued and refined by many others.[5-7]
Christe et al. introduced an (pseudo)isodesmic quantum chemical approach with
the experimental FIA of OCF2 of 209 kJ mol−1 as an anchor point, which avoids
the need to calculate a free fluoride anion (Equation 3.1).[8]
COF3− + A AF− + COF2
OCF3−COF2 + F−
A + F− AF−
DFT-Method
experimental
−ΔH = FIA
Equation 3.1: Christe’s definition of fluoride ion affinity (FIA).
By using this method the strength of Lewis acids have been ranked.[9] However a
downside of this classification is, that it only evaluates versus the, in the HSAB
sense,[10] hard fluoride ion. This works well for equally hard Lewis bases, but only
insufficient for HSAB soft Lewis bases. To estimate the strength of a Lewis acid
the affinities to ions of varying hardness (i.e., Cl−, H− and CH3− = Me−) should be
compared and the most suitable affinity value considered. Unfortunately, there is
no consistent reference system, which would allow comparing trends between
the different methods and lead to a broadly applicable Lewis acid scale.
However not only Lewis acids are in the focus of current interest, also weakly
coordinating anions (WCAs) have become an indispensable tool in chemistry and
found versatile applications (for more details see Chapter 1.5).[11, 12] In contrast to
a Lewis acid, not the strength of a WCA is in focus but its stability. Clearly the
stability of a typical WCA [M(L)n]− (M = Lewis Acidic central atom of valency n−1;
L = univalent residue) is related to the strength of the underlying Lewis acid
M(L)n−1 and thus is addressed in this work.
In independent work, the Krossing group estimated the relative stabilities of
WCAs based on calculations on a representative set of WCAs and their parent
Lewis acids. This method allowed to order the relative stabilities and coordinating
abilities of WCAs by the Ligand affinity (LA), decomposition in the presence of a
28
hard (proton decomposition PD) and a soft electrophile (copper decomposition
CuD), the position of the HOMO, the HOMO-LUMO gap and the before
mentioned FIA, which is capable to predict the stability of a given (WCA) by
viewing the FIA of its parent Lewis acid. The higher the FIA of the Lewis acid, the
more stable is the respective WCA towards ligand abstraction.
M(L)n− + L−
HL+ H+
+ Cu+ CuL−ΔH = CuD
M(L)n−
M(L)n−
ΔH = LA
−ΔH = PD
M(L)n−1
M(L)n−1
M(L)n−1
+
+
(a)
(b)
(c)
Equation 3.2: Definition of Ligand affinity (LA), proton decomposition (PD) and copper decomposition (CuD).
The relative ordering of the stabilities of the WCAs and their underlying Lewis
acids are reliable, because all calculations were done by the same method.
With this contribution we present such an approach augmented by hitherto only in
small parts addressed ion affinities in addition to the FIA. Thus, we compare the
Lewis acidity of the most common Lewis acids (A) with respect to their CIA
(Chloride Ion Affinity), HIA (Hydride Ion Affinity) and MIA (Methyl Ion Affinity). For
all three values and in addition also for the FIA, we chose a unified reference
system as anchor point. It is based on the trimethylsilyl compounds Me3SiY
(Y = F, Cl, H or Me) of the respective ions, so that the relative values of the Lewis
acidity towards different bases (Y−) are better comparable. A further advantage is
that only the respective reference reaction has to be optimized at a sufficiently
accurate and highly correlated level (here G3), while the residual in part very
large molecules can be assessed based on subsequent isodesmic reactions
calculated at a much less expensive level (here (RI-)BP86/SV(P)), Equation 3.3).
Me3SiY + A Me3Si+ + AY−
Me3Si+Me3SiY
A + Y− AY−
(RI-)BP86/SV(P)
G3 + Y−
Equation 3.3: Determining ion affinities with a unified reference system as anchor point.
29
To further validate the data we performed systematic calculations of the smaller
reference systems at the frozen core CCSD(T) level with correlation effects
extrapolated to a full quadruple-ζ basis.[13-15] The error bar of this methodology
was reported to be below 1 kJ mol−1,[14] and thus serves as a validation of the
simpler isodesmic procedure according to Equation 3.3 that for size reasons had
to be applied for larger Lewis acids.
Furthermore, we expanded the known scales for the WCA stability with the PD
and CuD of a series of hitherto not explored WCAs [M(L)n]−. All calculations were
done in the gas phase.
3.2 Results and Discussion
Background of these investigations have been experiments to activate hydrogen
with the Lewis acid tris(2H-hexafluoroisopropoxy)borane (B(Ohfip)3) 1, which
usually finds applications in electrochemistry.[16, 17] The fluorinated alkyl groups
increase with their electron-withdrawing group the anion complexation ability.[17]
As a consequence, the conductivity and lithium ion transference number in
lithium battery electrolytes is enhanced. We investigated the possibility to obtain
frustrated Lewis Pair (FLP) chemistry based on this Lewis acid and reacted it with
H2 and selected phosphanes, amines, pyridines and N-heterocyclic carbenes
(NHCs), but we have never observed a reaction. Even upon addition of very
strong Lewis bases, adduct formation was partly only occurring in equilibrium. By
contrast, the respective anion [B(Ohfip)4]− is known as a rather stable WCA and
the FIA of the B(Ohfip)3 acid, calculated according to the procedure in Equation
3.3, is with 384 kJ mol−1 rather large.[18, 19] Moreover, during the course of the
synthesis of [B(Ohfip)4]− from Na[BH4] and HO-hfip, the [H-B(Ohfip)3]− anion is a
stable intermediate that needs many hours of reflux to further completely react
with HO-hfip to give the symmetric borate. For both reasons, the high FIA of the
Lewis acid and the known stability of the [H-B(Ohfip)3]− anion, we did not expect
these unsuccessful reactions and therefore started to perform calculations to
investigate and compare the Lewis acidity towards HSAB-different ions for a
deeper insight. In the following we first describe our experiments, before turning
to the general calculations on a wide set of Lewis acids.
30
3.2.1 Synthesis and Characterization
We synthesized 1 according to the literature and obtained white plate-like crystals
melting at 32°C.[17, 20] The molecular structure of compound 1 was determined by
X-ray analysis.
Crystal Structures: The Lewis acid 1 crystallizes in the monoclinic space group
P21/c and exists, in contrast to the dimeric heavier homologue Al(Ohfip)3, as a
monomer.[21] The sum of the three O-B-O angles add up to 360°, so the orbital
interactions of the π-system between the boron and oxygen atoms are
maximized. The B-O distances of 136 pm are shortened by 11 pm compared to
the distances in the THF adduct of Na[B(Ohfip)]4.[18] The C-O distances are with
141 pm nearly unchanged compared to those of HO-hfip.[22] Interestingly, the
central B(OCH)3 unit resides in a plane. This might be attributed to i) the steric
demand of the hfip residues and ii) the formation of 3 weak intramolecular
(C-)H---O hydrogen bonds (dHO = 198 pm). Per ligand one CF3-group is above
and the other below this central plane.
31
O1
B1
O3
O2
C4
C1
C3
C2
C8
C5
C9
C6
C7 H2
H1
H3
F1
F2
F3 F4 F6
F5
F7
F9
F10
F8
F11F12
F13F14
F16F15 F17
F18
Figure 3.1: Molecular structure of 1. Thermal ellipsoids are shown at the 50% probability level, except fluorine atoms, which are reduced in size for clarity. Selected atom distances [pm] and angles [°]: B1-O1 = 136.0(4), B1-O2 = 135.8(4), B1-O3 = 135.9(4), O1-C1 = 141.6(3), O2-C2 = 141.6(3), O3-C3 = 141.1(3), O1-H1 = 198.1, O2-H2 = 198.0, O3-H3 = 197.8; O1-B1-O2 = 120.5(3), O2-B1-O3 = 119.7(3), O1-B1-O3 = 119.8(3), B1-O1-C1 = 122.3(2), B1-O2-C4 = 121.7(2), B1-O3-C7 = 122.3(2).
From a multitude of crystallization experiments with a great variety of neutral
Lewis bases, we obtained one very weak adduct by dissolving 1 in acetonitrile
(1•NCMe) and cooling the solution slowly down to −40°C.[23] The compound
1•NCMe crystallizes in the trigonal space group R3c. To the best of our
knowledge we report herein the longest coordination between boron and a
nitrogen atom of an acetonitrile molecule (248 pm). We were rather astonished to
see this weak interaction, as the boron nitrogen distances in related complexes
like MeCN•BCl3, MeCN•B(C6F5)3 or the acetonitrile adduct of the
perfluoroaryldiborane C6F4-1,2-(B(C6F5)2)2 are 156, 162 and 161 pm.[24, 25] The
boron atoms of these complexes are distorted tetrahedrally coordinated. The
hitherto longest boron nitrogen distance was measured in the mixed crystal
MeCN• (B(CH2Ph)3)0.92(Ga(CH2Ph)3)0.08 and amounts to 178 pm.[26]
32
F1F2
F3
F4
F6F5
F7
F9
F10
F8
F11
F12
F13F14
F16
F15
F17
F18B1
O3
O2C4
C1
C3
C2
C8
C5
C9C6
C7
H2H1
H3
H6H4
C11
N1
C10
O1
H5
Figure 3.2: Molecular structure of 1•NCMe. Thermal ellipsoids are shown at the 50% probability level, except fluorine atoms, which are reduced in size for clarity. Selected atom distances [pm] and angles [°]: B1-N1 = 248,4, B1-O1 = 136.97(9), O1-C1 = 140.46(16), O1-H1 = 197.0, N1-H1 = 299.4; O1-B1-O2 = 119.57(3), B1-O1-C1 =121.31(9).
The O-B-O angles in 1•NCMe amount on average to 119.6° and add up to
358.7°, which signals an almost ideal trigonal planar geometry. The averaged B-
O bond length is 137 pm and hence only very slightly widened compared to 1
(136 pm). The acetonitrile molecule of 1•NCMe lies on a pseudo C3 axis. On a
nitrogen atom of one acetonitrile molecule follows the methylgroup of a second.
Between these molecules are the boron atoms of B(Ohfip)3, whereupon the
distance from the boron atom to the carbon atom of the methylgroup of a second
MeCN•B(Ohfip)3 adduct is 421 pm.
FLP Chemistry: We continued our investigation by testing the ability of 1 to
activate hydrogen with different FLP compounds. To analyze the
thermodynamics of the hydrogen activation of 1 in comparison to the well-known
B(C6F5)3 the standard Gibbs energies were calculated by DFT calculations
((RI-)BP86/SV(P)[27, 28] with the COSMO solvation model[29]) in the gas phase and
different solvents by using a Born-Fajans-Haber cycle. In agreement with our
futile experimental efforts, FLP chemistry based on 1 is considerably less favored
33
than the one with B(C6F5)3. Furthermore, these investigations suggested that
polar solvents promote H2 activation and therefore acetonitrile was inter alia
selected as a solvent. There is a trend towards stronger Lewis bases observable,
so, because of their high steric demand and convenient energy levels, NHC
based compounds should be suitable bases.
Lews acid(g) Lewis base(g) H2(g) [H-Lewis acid]−(g) [H-Lewis base]+(g)
Lewis acid(solv) Lewis base(solv) H2(solv) [H-Lewis acid]−(solv) [H-Lewis base]+(solv)
−ΔsolvG° −ΔsolvG° −ΔsolvG° −ΔsolvG° −ΔsolvG°
−ΔRG°(g)
−ΔRG°(solvent)
Scheme 3.1: Born-Haber-Fajans cycle to access ΔRG0(g) and ΔRG0
(solvent) of 1 and B(C6F5)3 in various solvents.
34
Table 3.1: ΔRG0(g) and ΔRG0
(solvent) of the H2-activation of selected Lewis acid/base-pairs in various solvents ((RI-)BP86/SV(P)) and calculated according to the cycle in Scheme 3.1; NHC = 1,3-dimethyl-4,5-diphenyl imidazole-2-ylidene, IDipp = 1,3-bis(2,6-diisopropylphenyl)imidazol-2-yliden.
Compound ΔRG0(g) ΔRG0(CH2Cl2) ΔRG0(o-difluorobenzene) ΔRG0(H3CCN)
1
PPh3 464 206 188 170
PtBu3 419 146 128 109
NEt3 475 162 141 118
2,6-lutidine 473 169 149 127
NHPh2 552 250 229 207
NHC 310 51 34 15
IDipp 309 67 51 34
B(C6F5)3
PPh3 292 55 39 22
PtBu3 247 −4 −21 −39
NEt3 303 11 −8 −29
2,6-lutidine 301 18 −1 −21
NHPh2 380 99 80 59
NHC 138 −100 −116 −133
IDipp 137 −84 −98 −114
We started an investigation with several common Lewis bases. All following
reactions were carried out under an argon atmosphere in NMR tubes with a J.
Young valve according to the following procedure: Compound 1 and the different
phosphanes, amines and NHCs were mixed in a 1:1 stoichiometry in
d3-acetonitrile and characterized by NMR spectroscopy. Subsequently the
reaction mixtures were evacuated, exposed to hydrogen pressure of one bar, and
were then again NMR spectroscopically analyzed. We started our attempts by
using phosphanes, since e.g. B(C6F5)3 i) forms adducts with PPh3 or Cp2PH,[25, 30]
ii) in combination with (C6F5)Ph2P it leads to a FLP system capable to reversible
activate hydrogen[31] and iii) a mixture with PtBu3 is able to activate terminal
alkynes.[32] Here PPh3, PtBu3 and PMePh2 did not form classical adducts with 1,
and there is no sign of a hydrogen activation (e.g. doublet 1JPH in the 31P NMR,
doublet 1JBH splitting in the 11B NMR spectrum: expected for [HB(Ohfip)3]− 127 Hz
at δ = 7.7 ppm). To eliminate the possibility that the reaction occurs very slowly,
35
in a second approach 1 was stirred with PPh3 over twenty days and exposed to
H2 at 1 bar, but without any noticeable reaction. Of all the tested bases collected
in Table 3.1, 1 only forms an adduct visible in the NMR with NEt3. However, this
adduct only exists in equilibrium. We tried to obtain crystals of the adduct, but we
only obtained crystals of 1•NCMe. Also testing the ability to activate hydrogen
failed in all instances. No adduct formation and no hydrogen cleavage were
detected by the combination of 1 and NHPh2 or 2,6-lutidin, which is able to
activate hydrogen in combination with B(C6F5)3.[33] Since it appeared that the
reactivity of 1 is not sufficient to activate H2 with typical Lewis bases like
phosphanes and amines, we used the strongly basic NHCs 1,3-dimethyl-4,5-
diphenyl imidazol-2-ylidene and 1,3-di(2,6-diisopropylphenyl) imidazol-2-ylidene.
They form classical adducts with 1, which were observed in the 11B NMR spectra
at 1.5 and accordingly 1.9 ppm, but no detectable hydrogen activation occurred.
NN
P NP
P
N NH NN
Figure 3.3: Tested Lewis basis for hydrogen activation with B(Ohfip)3.
Furthermore we investigated, if the Li+ salt of the [HB(Ohfip)3]− anion and a
protonated Lewis base like [HNEt3]+Cl− are compatible in solution or if they
release elemental hydrogen. For these screening experiments, crude
Li+[HB(Ohfip)3]− contaminated with about 75 mol-% Li[Ohfip] was used. We
dissolved [HNEt3]+Cl− and the contaminated Li[HB(Ohfip)3] salt in acetonitrile to
preliminary test their compatibility. However, even after stirring for one week no
gas formation occurred and the NMR signals of the starting materials remained
unchanged. Thus, no regeneration of the original Lewis acid/Base pairs occurred.
Therefore, we have to note that H2 activation might still be possible with this
system, but at least not with our tested reaction conditions.
36
3.2.2 Quantum Chemical Investigations
Ion Affinities of Lewis acids A: Earlier and the following calculations to
determine ion affinities are based on isodesmic reactions with the BP86
functional[27] and the SV(P) basis,[28] which represent a good agreement between
costs and accuracy. This method allows fast access to large molecules with up to
100 atoms, which cannot be calculated at correlated levels. During our
investigation to improve these scales by unifying the reference system and
adding up the CIA, HIA and MIA values, we noticed a larger discrepancy to
earlier published FIA values for BI3, AlCl3 and AlI3.[9]
Validation Study: To validate and confirm the BP86/SV(P) values, we carried
out RI-MP2 structure optimizations with TURBOMOLE[34] and def2-QZVPP basis
sets[35] and corresponding RI-C auxiliary bases for all atoms.[36] Based on these
structures, gas phase reaction energies were calculated with single point
calculations at the CCSD(T)(FC)/double-ζ level plus an MP2 extrapolation of the
correlation energy from double-ζ to quadruple-ζ basis sets with Gaussian 03
(Table 3.2). This approach was published earlier by Klopper et al.[13, 14] and has
successfully been used by our group to study protonation equilibria.[3]
A + Y− AY−−ΔH = YIA
Y = Cl ,H ,F
Equation 3.4: General equation to determine the ion affinity of different ions.
ΔEextrapol = ΔECCSD(T)/AVDZ − ΔEMP2/AVDZ + ΔEMP2/AVQZ
37
Table 3.2: Overview of frozen core and basis sets for all used elements.
element fc for MP2/def2-QZVPP AVXZ (X = D, Q) fc for AVXZ
H 0 aug-cc-pVXZ[37] 0
B 1 aug-cc-pVXZ[37] 1
C 1 aug-cc-pVXZ[37] 1
N 1 aug-cc-pVXZ[37] 1
O 1 aug-cc-pVXZ[37] 1
F 1 aug-cc-pVXZ[37] 1
Al 5 aug-cc-pV(X+d)Z[38] 5
P 5 aug-cc-pV(X+d)Z[38] 5
Cl 5 aug-cc-pV(X+d)Z[38] 5
Ga 14 aug-cc-pwCVXZ-PP[a][39, 40] 0
As 9 aug-cc-pwCVXZ-PP[a][39, 40] 0
Br 14 aug-cc-pVXZ-PP[41] 9
Sb 4 aug-cc-pwCVXZ-PP[a][39, 40] 0
[a] Constructed by combining cc-pWCVQZ-PP[40] with aug-cc-pVQZ-PP Diffuse.[39]
Thermal contributions to the enthalpy were calculated with BP86/def-TZVP at 1
bar, 298.15 K.
ΔU0 = ΔEextrapol + ΔEvrt
(ΔEvrt = sum of translation, rotational, vibrational energy inclusive zero point
vibrational energy)
ΔH0 = ΔU0 + RT
38
Table 3.3: Overview on the CIA, HIA and FIA values [kJ mol−1] of the chosen representative Lewis acid set calculated via CCSD(T)(FC)/double-ζ level plus MP2 extrapolation to quadruple-ζ. The values in parentheses give the difference to the affinity values calculated via Equation 3.3 at the BP86/SV(P) level. [SbF5H]− is not stable and would decompose to SbF4
− + HF, this value is given in square brackets.
Lewis acid CIA HIA FIA
BF3 151 (5) 297 (−2) 346 (4)
BCl3 195 (12) 395 (5) 384 (8)
BBr3 219 (7) 440 (2) 425 (−16)
AlF3 308 (1) 388 (−36) 482 (11)
AlCl3 320 (3) 428 (−22) 502 (4)
AlBr3 324 (−2) 425 (−39) 505 (−5)
GaF3 319 (14) 444 (−18) 447 (13)
GaCl3 299 (5) 446 (−19) 429 (−5)
GaBr3 295 (1) 445 (−24) 426 (−13)
PF5 165 (−14) 400 (−17) 380 (−17)
PCl5 179 (2) 468 (−14) 393 (1)
AsF5 237 (−14) 461 (−24) 434 (4)
SbF5 333 (−8) 530 [517] (−32/−13) 495 (2)
B(CN)3 363 (12) 583 (−3) 540 (−10)
B(OH)3 a) 163 (−7) 208 (4)
a) Does not form complexes.
The CIA and FIA values in Table 3.3 agree well with the values calculated on
BP86/SV(P), with the largest deviations being up to 17 kJ mol−1. The BP86/SV(P)
HIA values of the boron halides, B(CN)3 and B(OH)3 nicely fit, since the B-H bond
is little polar. Nevertheless, if the central atom is not a second row element and
the E-H bonds get more polar or hydridic, the HIA values are roughly 20 kJ mol−1
too high (maximum deviation: 39 kJ mol−1 for AlBr3). In order to obtain better
values, we also calculated the HIA values with BP86/SVP and thus p-type
polarization functions also at the hydrogen atom, but those values were even
inferior to those calculated with BP86/SV(P). Also orienting BP86/TZVP
calculations were inferior to the simple BP86/SV(P) method. Thus, for simplicity
and to be applicable for a larger set of compounds, we used the BP86/SV(P)
method and the procedure described in Equation 3.3. Since the difference in the
HIA values from the absolute values always points to one direction, the relative
HIA values at the simpler BP86/SV(P) level should still be suitable for a
discussion.
39
Ion Affinity Scale: With this cadre of validated data, we calculated the FIA, CIA,
HIA and MIA through a set of isodesmic reactions. Basis reaction of the
respective ion affinity calculations were the ion affinities of Me3SiY (Y = F, Cl, H,
Me), which were calculated at the reliable G3 level (Equation 3.3 and Table 3.4).
Table 3.4: Calculated reference reaction values (G3 level) as anchor points to determine FIA, CIA, HIA and MIA values.
Reference system ΔH [kJ mol−1]
Me3SiF → Me3Si+ + F− ±958
Me3SiCl → Me3Si+ + Cl− ±759
Me3SiH → Me3Si+ + H− ±959
Me3SiMe → Me3Si+ + Me− ±1000
Halide Acids: The HIA and MIA values of main group III halides (Table 3.5)
follow similar trends and rise for the heavier halogen atoms. If we take a look at
the affinity values of AlF3 and GaF3 there is not such a large difference compared
to the difference between AlF3 and BF3, since aluminum and gallium have nearly
the same size and aluminum is just a little bit more electro positive than gallium.
However, they distinguish in their affinity values towards soft or hard Lewis
bases. AlF3 favors the hard fluoride ion, while GaF3 prefers softer Lewis bases
like Me− or H− anions, which confirms that aluminumtrifluoride is harder than
galliumtrifluoride. Nevertheless, if the halides ligands get heavier, gallium turns
out scaled by aluminum, which indicates that the GaX3 Lewis acids are more
stable and less acidic. This could mean that the overlap of the p orbitals is more
pronounced and hence the π-back bonding is stronger. Summarized, in this row,
the AlX3 compounds are the strongest Lewis acids with respect to fluoride and
chloride, but towards the hydride ion, GaX3 acids are stronger and with respect to
the methanide ion similar to the aluminum acids.
The CIA and FIA values of PF5 and PCl5 are quite low and nearly identical, which
may be a result of the electrostatic repulsion of the six rather hard halides in the
octahedral complexes. For the strongest classical Lewis acid SbF5, two HIA and
MIA values are given, since [SbF5H]− and [SbF5Me]− are not stable and would
decompose.
40
41
Putting Boron Acids in Context: As expected B(OH)3 is a quite weak Lewis
acid and since [B(OH)3Cl]− is experimentally unknown we disregarded this anion
from our calculations. Compound 1 reveals relatively low ion affinity values.
Compared to B(C6F5)3 the FIA value of 1 is just 68 kJ mol−1 lower, but the HIA
value is 136 kJ mol−1 lower. This huge difference in the HIA value may explain
our experimentally observed results.
Interestingly the HIA of B(C6F5)3 is nearly identical with that one of the versus
fluoride stronger Lewis Acid Al(C6F5)3. B(C12F9)3 and B(C6H3-3,5-(CF3)2)3 show
the same effect, which implies that boron Lewis acids stabilize the hydride ion
better than other ions in the corresponding anion. The low CIA and FIA, MIA
values of B(C12F9)3 results from the sterically demanding perfluorinated biphenyl
ligands, which hinder the access to the Lewis acid center. Since F4C6(1,2-
(B(C6F5)2)2 has two nearby Lewis acid atoms small ions like fluoride and hydride
prefer a bridged geometry, but the difference between the bridged and not
bridged geometry is smaller than 1 kJ mol−1. PF2(C2F5)3 possesses compared to
PF5 slightly lower CIA, FIA and MIA values, due to the steric demand of the C2F5
groups in the acid base complexes. The HIA of PF2(C2F5)3 is a bit higher, since
H− is small and the steric effect of the C2F5 groups is overcompensated by its
electron withdrawing effect. Main group IV is represented by the silylium
zwitterion SiMe2CH2CB11Cl11[42] and the Janus-headed Lewis acid TMS-F-
Al(OC(CF3)3)3.[43] The ion affinity of the silicon ion is boosted by around 140 kJ
mol−1 compared to the silicon atom in neutral TMS-F-Al(OC(CF3)3)3.
Lewis Superacids: Since SbF5 is viewed as the strongest conventional Lewis
acid, the FIA value of monomeric SbF5 is used to determine Lewis superacids:
“Molecular Lewis acids, which are stronger than monomeric SbF5 in the gas
phase are Lewis superacids.”[9] This approach is simple and since strong Lewis
acids and the fluoride anion are typically hard, it gives a nice overview of Lewis
acids by their relative strengths. The FIA values of monomeric AlCl3, AlBr3, AlI3,
Al(OC(CF3)3)3, B(OTeF5)5, As(OTeF5)5, Sb(OTeF5)5, B(CN)3, B(CF3)3, Al(C6F5)3,
F4C6(1,2-(B(C6F5)2)2 and SiMe2CH2CB11Cl11 excel the FIA value of SbF5 and may
be classified as Lewis superacids.
Table 3.5: Overview on the calculated CIA, HIA, FIA and MIA values of representative Lewis Acids. [SbF5H]− and [SbF5Me]− are less stable and would decompose to SbF4− + HF
and SbF4− + MeF, respectively, this value is given in brackets. TMS-F-Al(OC(CF3)3)3 can be attacked, depending on the nucleophile, at the Si or the Al atom and decomposes either
to TMS-Y + F-Al(OC(CF3)3)3 or TMS-F + Y-Al(OC(CF3)3)3. The value for the nucleophilic attack at the Al atom given in parenthesis. [a] The anion containing a B-Y-B bridge is thermodynamically favored. [b] The most stable isomer was used. [c] The experimental crystal structure was used as start geometry.[44] Non-isolable acids and their ion affinity values are given in italics.
Lewis Acid LUMO [eV]
CIA [kJ mol−1]
HIA [kJ mol−1]
FIA [kJ mol−1]
MIA [kJ mol−1]
Lewis Acid LUMO [eV]
CIA [kJ mol−1]
HIA [kJ mol−1]
FIA [kJ mol−1]
MIA [kJ mol−1]
M-X B(O 3 hfip) −0.42 141 348 384 387 BF3 −0.29 146 299 342 355 Al(OC(CF3)3)3 −1.51 352 490 543 530 BCl3 −2.38 183 391 405 436 B(OTeF5)3 −6.78 325 556 552 602 BBr3 −2.88 213 438 441 477 As(OTeF5)5 −7.30 403 710 559 753 BI3 −3.41 261 505 493 540 Sb(OTeF5)5 −7.62 465 746 625 809 AlF3 −2.24 306 423 471 464 M-C AlCl3 −2.18 318 450 498 490 B(CN)3 −5.97 351 587 551 610 AlBr3 −2.48 326 464 510 502 B(CF3)3 −4.77 358 583 556 614 AlI3 −3.01 347 497 535 535 B(C6F5)3 −3.93 236 484 452 483 GaF3 −3.41 306 462 434 491 Al(C6F5)3 −3.07 348 483 536 518 GaCl3 −3.13 294 464 434 493 Ga(C6F5)3 −3,28 307 479 453 502 GaBr3 −3.34 295 470 438 498 B(C12F9)3 −3.95 190 452 431 415 GaI3 −3.74 310 495 457 523 B(C6H3-3,5-(CF3)2)3 −4.04 281 486 482 504 PF5 −1.53 179 417 398 456 B(C10F7)3 −3.91 265 519 483 519 PCl5 −4.73 178 483 392 507 F4C6(1,2-(B(C6F5)2)2 −4.11 309 536[a] 523 493[a]
AsF5 −4.22 251 485 430 527 B2(C6F5)2(C6F4)2 −4.68 271 514 477 524 SbF5 −5.83 341 562 [531] 493 607 [544] PF2(C2F5)3 −2.42 157[b] 428[b] 388[c] 426[b] M-O SiMe2CH2CB11Cl11 −3.03 397 590 597 627 B(OH)3 −0.13 170 204 220 TMS-F-Al(OC(CF3)3)3 −1.21 259 (267) 459 (407) 458 459 (407)
42
Stability of WCAs based on FIA, PD, CuD, HOMO level and HOMO-LUMO gap: If the before mentioned Lewis acids are expanded by a L− ligand, the
related WCA is obtained. We determined the stability towards decomposition,
oxidation and reduction for most of the WCAs that relate to the Lewis acids in
Table 3.5.
The higher the FIA of the acid the more stable is the WCA towards ligand
abstraction. To rate the stability of a WCA towards attack of a hard (H+) and a
soft electrophile (Cu+) the isodesmic decomposition reactions (eq. 2 and 3) were
calculated to obtain the proton decomposition (PD) and the copper
decomposition (CuD). Herein we show instead of the previously used ΔrU values,
the ΔrG° values of the PD and CuD, hence they are closer to laboratory
conditions. The entropy S of the H+ and Cu+ cations was calculated by using the
Sackur-Tetrode equation.[45] Since a gaseous anion and a gaseous cation react
to give two neutral species, the PD and CuD are both exergonic. The less
negative the PD and CuD values are, the more stable is the WCA against
electrophilic attack.[7] The lower the HOMO energy the more resistant is an anion
towards oxidation, hence the electron is harder to remove. The HOMO-LUMO
gap is related to its resistance towards reduction. The larger the gap, the more
stable is the WCA towards gaining an electron. The data in Table 3.5 cannot be
taken as absolute, but since the same calculation methods were used, relative
trends will definitely be correct. However we have to keep in mind that just the
most frequently experimental observed decomposition pathways are calculated
and of course other decomposition pathways may occur.
43
Table 3.6: Calculated properties of WCAs: FIA of the parent Lewis acid as shown in Table 3.5.
Anion FIA [kJ mol−1]
PD [kJ mol−1]
CuD [kJ mol−1]
HOMO [eV]
Gap [eV]
M-X [BF4]− 342 −1212 −540 −1.799 10.820[BCl4]− 405 −1237 −629 −1.708 7.975[BBr4]− 441 −1225 −631 −1.787 5.820[BI4]− 493 −1190 −613 −2.127 3.620[AlF4]− 471 −1083 −411 −2,510 8.016[AlCl4]− 498 −1102 −493 −2.546 7.054[AlBr4]− 510 −1105 −511 −2.505 5.685[AlI4]− 535 −1095 −519 −2.658 3.973[GaF4]− 434 −1118 −446 −2.677 7.048[GaCl4]− 434 −1125 −516 −2.604 5.501[GaBr4]− 438 −1127 −533 −2.532 4.369[GaI4]− 457 −1112 −536 −2.673 3.051[PF6]− 398 −1163 −491 −2.673 8.801[PCl6]− 392 −1242 −633 −2.246 1.929[AsF6]− 430 −1129 −457 −3.150 6.282[SbF6]− 493 −1065 −393 −3.911 5.134M-O [B(Ohfip)4]− 384 −1212 −526 −3.377 7.056[Al(OC(CF3)3)4]− 543 −1077 −413 −4.096 6.737[B(OTeF5)4]− 552 −1098 −496 −5.547 2.126[As(OTeF5)6]− 559 −1053 −452 −6.129 1.983[Sb(OTeF5)6]− 625 −999 −398 −6.460 2.181M-C [B(CN)4]− 551 −1092 −438 −4.182 6.818[B(CF3)4]− 556 −1143 −411 −3.527 9.069[B(C6F5)4]− 452 −1263 −567 −3.120 4.214[Al(C6F5)4]− 536 −1224 −528 −3.304 4.251[Ga(C6F5)4]− 453 −1246 −550 −3.308 4.332[B(C12F9)4]− 431 −1231 −534 −3.517 3.339[B(C6H3(CF3)2)4]− 482 −1250 −527 −3.798 3.930[B(C10F7)4]− 483 −1236 −540 −3.098 2.795[F4C6(1,2-(B(C6F5)2)2)(C6F5)]− 523 −1325 −629 −3.207 1.969[B2(C6F5)3(C6F4)2]− 477 −1259 −563 −3.284 2.548
For the [MX4]− anions (M = B, Al, Ga; X = F, Cl ,Br, I) it can be seen from Table
3.6, that the [MF4]− and [MI4]− anions have an increased resistance against an
attack of a soft or a hard nucleophile (see PD and CuD). The [MF4]− and [MCl4]−
anions are exceptional stable towards reduction (see gap), which is even more
pronounced in the phosphor halides. This confirms that especially fluorine is a
suitable ligand at the central atom to enhance the properties of a WCA. Usually
the HOMO level rises with increasing weight of the central atoms and ligands.
44
45
The alkoxyaluminate [Al(OC(CF3)3)4]− possesses distinguished thermodynamic
stability values paired with a simple straightforward synthesis. Compared to the
teflate based anions of Table 3.6 it offers similar values (see FIA, PD, CuD, and
HOMO level), but it is significant more reduction resistant. Additionally it is not
very moisture sensitive and can be easily handled in glass. These values
underline its current role in chemistry to stabilize highly reactive cations.[6, 12, 46]
Furthermore, Table 3.6 includes a series of WCAs with mainly fluorinated organic
ligands. Of those [B(CN)4]− and [B(CF3)4]− offer very good WCA properties,
although [B(CF3)4]− allow for additional decomposition pathways.[47] Typically, the
fluorination of ligands increases the stability values of borate based anions.[7]
Apart from that, the stability values of all the M-C bonds are quite similar.
3.2.3 Analysis
Relationship between Ion Affinity Values and LUMO level: To gain a deeper
insight in the relationship between the ion affinity values and the LUMO levels of
our spotlighted Lewis acids of Table 3.5, we plotted the values and added a
regression line in each case (see Figure 3.4).
-8 -7 -6 -5 -4 -3 -2 -1 0100
150
200
250
300
350
400
450
500
CIA
[kJ/
mol
]
LUMO [eV]
-8 -7 -6 -5 -4 -3 -2 -1 0
200
300
400
500
600
700
800
900
MIA
[kJ/
mol
]
LUMO [eV]
-8 -7 -6 -5 -4 -3 -2 -1 0100
200
300
400
500
600
700
800
HIA
[kJ/
mol
]
LUMO [eV]
a) b)
c) d)
-8 -7 -6 -5 -4 -3 -2 -1 0
200
300
400
500
600
700
FIA
[kJ/
mol
]
LUMO [eV]
Figure 3.4: Plots showing the regression lines of the relationships between the ion affinity values and the LUMO level of the Lewis acids of Table 3.5. Linear equation of ion affinity values against the LUMO level for a): y = −30.14(6.05)x + 364.61(24.86), R2 = 0.41, b): y = −26.31(6.71)x + 189.68(26.85), R2 = 0.31, c): y = −48.68(5.06)x + 311.86(19.98), R2 = 0.73 and d): y = −46.90(5.75)x + 345.77(22.68), R2 = 0.67.
By trend the gradient shows, that a low LUMO level is accompanied with a high
ion affinity value, however a low LUMO level is not automatically connected with
a high ion affinity value. Furthermore R2 and the value of the gradient of the FIA
and CIA values are compared to the corresponding values of the HIA and MIA
values significantly smaller. These facts are consistent with Pearsons HSAB
concept, since interaction of the soft hydride and methyl ions with Lewis bases
are more orbital based and consequently strongly influenced by the LUMO levels.
Therefore, the HIA and MIA values are in a distinctive relationship with the LUMO
values. Especially the HIA can be well predicted by the LUMO level of a given
46
Lewis acid and vice versa, since there is no steric influence on the result. On the
other hand, the interaction of the hard fluoride and chloride ions has a high ionic
contribution and thus it is less attached to the LUMO energies. On first sight, the
FIA should have an even lower R2 value than the CIA, because the fluoride ion is
harder than the chloride ion, but the chemistry of fluoride is sometimes
exceptional. Since the fluoride atom is smaller than the chloride, it is possible that
the overlapping of the involved orbitals is improved and steric effects are less
developed. Therefore, it gets a more covalent character and fits the trend better
than chloride.
Surface Analysis of 1•NCMe: The long B---N distance in the crystal structure of
1•NCMe (Figure 3.2) gives reason for a more detailed look. The red color in the
Hirshfeld surface[48] (Figure 3.5 left) indicates a distance below the sum of van
der Waals radii. Additionally we performed calculations of the electrostatic
potential on an electron density surface with varied isovalues (0.004, 0.016,
0.018, 0.019, 0.020, 0.021, 0.030, 0.040, 0.060 and 0.080 e au−3), but no
significant B-N interaction was recognized in one of those layers.[49] Figure 3.5
right shows that intramolecular interactions are more pronounced than the
intermolecular interactions between boron and nitrogen (weak polarization).
Furthermore, there is no sign of an N-H or N-F interaction.
Figure 3.5: Left: Hirshfeld surface of 1●NCMe. Right: Electrostatic potential on an electro density surface of 1●NCMe (isovalue = 0.040 e au−3).
47
3.3 Conclusion
Herein we introduced the so far unknown crystal structure of B(Ohfip)3 and
discovered an extraordinary long B---N distance in the crystal structure of the
acetonitrile adduct. Attempts to use B(Ohfip)3 as Lewis acid component in FLP
chemistry systems failed at the used reaction conditions. To investigate this result
we introduced a new validated Lewis acid scale based on a consistent reference
system (MeSiY; Y = Cl−, H−, F−, Me−) and calculated the CIA, HIA, FIA and MIA
values of a multitude of common and frequently used Lewis acids (Table 3.5).
Lewis acidity depends on many different factors, like the LUMO level, stabilizing
interactions in the Lewis acid, electro-negativity and steric demand of the ligands,
reorganizing energy, solvent effects etc. and especially on the bonding partner,
the Lewis base. With our scale, it is possible, given that the reference
calculations are done, to obtain comparable ion affinity values to four different
Lewis bases of any given Lewis acid by performing only five low level
calculations. Additionally we evaluated the stability of WCAs based on the Lewis
acids by calculating the LUMO energies, HOMO-LUMO gap, proton
decomposition (PD) and the copper decomposition (CuD).
3.4 Experimental Section
3.4.1 Theoretical Methods
All geometries were optimized at the (RI-)BP86/SV(P) level[27, 28] using
TURBOMOLE 6.4.[34] Vibrational frequencies were calculated with the AOFORCE
module[50] and all structures represented true minima without imaginary
frequencies. Solvation enthalpies in dichloromethane, ortho-difluorobenzene and
acetonitrile were calculated with the COSMO module[29] using the default options
and standard optimized COSMO radii (single points on optimized structures).
Permittivity values (CH2Cl2: εr = 8.93, o-difluorobenzene: εr = 14.3, H3CCN: εr = 36.6) were taken from the literature.[51] Highly correlated calculations were
done according to the method of Klopper et al.[13, 14] RI-MP2 structure re-
optimizations were carried out with TURBOMOLE 6.4[34] and def2-QZVPP basis
sets[35] and corresponding RI-C auxiliary bases[52] for all atoms. MP2 and ccsd(t)
48
single point calculations with correlation-consistent basis sets were done with
Gaussian 09.[53] The reference values for the Lewis acidity scale were calculated
with the G3 method.[54]
3.4.2 Synthesis and Characterization
Techniques and Instruments: All reactions were carried out under an inert
atmosphere by using standard vacuum and Schlenk techniques or a glovebox
with an argon atmosphere (H2O and O2 < 1 ppm), Special J. Young NMR tubes
sealed with Teflon valves were used to exclude air and moisture. All solvents
were dried over CaH2 or P4O10 and distilled afterwards. NMR data were recorded
from solutions in d8-toluene or d3-acetontrile at room temperature or 313 K on a
BRUKER AVANCE II+ 400MHz WB spectrometer. 1H and 13C chemical shifts are
given with respect to TMS, 19F NMR spectra to fluorotrichloromethane, 11B NMR
spectra to the boron-trifluoride-diethyl-ether-complex, 31P NMR spectra to a 85 %
phosphoric(V) acid solution and 7Li NMR spectra to 9.7 M LiCl in D2O. IR spectra
were recorded on a Diamond ATR unit (200-30000 cm−1) on a BRUKER alpha
Fourier transform-infrared (FT-IR) spectrometer with a resolution of 4 cm-1.
Raman spectra were obtained on a BRUKER VERTEX 70 spectrometer with a
BRUKER RAM II Raman unit in sealed melting-point capillaries. Data collections
for X-ray structure determinations were performed on a Rigaku Spider image
plate system or a BRUKER APEX II Quazar CCD diffractometer at 100 and 110
K, respectively, with Mo radiation. The single crystals were mounted in
perfluoroether oil on a MiTeGen MicromountTM. 1,3-dimethyl-4,5-diphenyl
imidazol-2-ylidene was synthesized from the thiourea precursor according to the
literature.[55]
Preparation of compound 1: Compound 1 was synthesized according to the
literature,[17] but the reaction was executed in CH2Cl2 instead of pentane and
further purified by sublimation and crystals suitable for X-ray diffraction were
obtained. 1H NMR (400.17 MHz, CD3CN): δ = 1.96 (quint., 1H, CHD2CN), 5.48
(sept., 3H, B(Ohfip)3). 19F NMR (376.54 MHz, CD3CN): δ = −75.8 (d, 18F,
B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN): δ = 17.3 (br, 1B, B(Ohfip)3). IR
(Diamond ATR, corrected): ν~= 650 (cw), 690 (w), 745 (w), 874 (w), 908 (w), 1078
49
(w), 1109 (vs), 1200 (s), 1230 (m), 1279 (s), 1383 (m), 1406 (m), 1454 (w), 2985
(vw) cm−1. Raman: ν~ = 210, 277, 299, 335, 367, 391, 431, 543, 568, 649, 692,
731, 746, 860, 902, 1112, 1209, 1253, 1280, 1309, 1385, 2739, 2757, 2937,
2985 cm−1.
Preparation of compound 1•NMe: The crystals of 1•NMe were obtained by
dissolving 1 in acetonitrile, cooling the solution slowly down to −40°C and
keeping these conditions for two weeks.
Analysis of the system 1/PPh3: Compound 1 (165 mg, 0.3 mmol, 1 eq.) and
PPh3 (85 mg, 0.3 mmol, 1 eq.) were mixed in a J. Young NMR tube and
dissolved in d3-acetonitrile. The reaction mixture was characterized by NMR
spectroscopy. Subsequently the reaction mixture was evacuated, exposed to
hydrogen pressure of one bar, and was then again NMR spectroscopically
analyzed.
In a second approach an identically prepared reaction mixture was stirred 20
days, NMR spectroscopially characterized, evacuated, and exposed to hydrogen
pressure of one bar for 30 minutes and NMR spectroscopically analyzed. All four
different solutions (before and after adding hydrogen) were NMR
spectroscopically identical. 1H NMR (400.17 MHz, CD3CN): δ = 1.96 (quint., 1H,
CHD2CN), 5.48 (sept., 3H, B(Ohfip)3), 7.38 (m, 15H, PPh3).19F NMR (376.54
MHz, CD3CN): δ = −75.8 (d, 18F, B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN):
δ = 17.3 (br, 1B, B(Ohfip)3). 31P NMR (376.54 MHz, CD3CN): δ = −5.4 (m, 1P,
PPh3).
Analysis of the system 1/PtBu3: Compound 1 (165 mg, 0.3 mmol, 1 eq.) and
PtBu3 (63 mg, 0.3 mmol, 1 eq.) were mixed in a J. Young NMR tube and
dissolved in d3-acetonitrile. The reaction mixture was characterized by NMR
spectroscopy. Subsequently the reaction mixture was evacuated, exposed to
hydrogen pressure of one bar, and was then again NMR spectroscopically
analyzed. The two solutions (before and after adding hydrogen) were NMR
spectroscopically identical. 1H NMR (400.17 MHz, CD3CN): δ = 1.32 (d, 27H,
PtBu3), 1.96 (quint., 1H, CHD2CN), 5.48 (sept., 3H, B(Ohfip)3).19F NMR (376.54
MHz, CD3CN): δ = −75.8 (d, 18F, B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN):
50
δ = 17.2 (br, 1B, B(Ohfip)3). 31P NMR (376.54 MHz, CD3CN): δ = 62 (m, 1P,
PtBu3).
Analysis of the system 1/PMePh2: Compound 1 (131 mg, 0.25 mmol, 1 eq.)
and PMePh2 (50 mg, 0.25 mmol, 1 eq.) were mixed in a J. Young NMR tube and
dissolved in d3-acetonitrile. The reaction mixture was characterized by NMR
spectroscopy. Subsequently the reaction mixture was evacuated, exposed to
hydrogen pressure of one bar, and was then again NMR spectroscopically
analyzed. The two solutions (before and after adding hydrogen) were NMR
spectroscopically identical. 1H NMR (400.17 MHz, CD3CN): δ = 1.64 (s, 3H,
PMePh2), 1.96 (quint., 1H, CHD2CN), 5.48 (sept., 3H, B(Ohfip)3), 7.38 (m, 6H,
PMe(o/p)Ph2), 7.45 (m, 4H PMe(m)Ph2). 19F NMR (376.54 MHz, CD3CN):
δ = −75.8 (d, 18F, B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN): δ = 17.3 (br, 1B,
B(Ohfip)3). 31P NMR (376.54 MHz, CD3CN): δ = 27 (br, 1P, PMePh2).
Analysis of the system 1/NEt3: Compound 1 (250 mg, 0.5 mmol, 1 eq.) and
NEt3 (50 mg, 0.5 mmol, 1 eq.) were mixed in a J. Young NMR tube. The reaction
mixture turned immediately yellow. Afterwards the reaction mixture was dissolved
in d3-acetonitrile and characterized by NMR spectroscopy. Subsequently the
reaction mixture was evacuated, exposed to hydrogen pressure of one bar, and
was then again NMR spectroscopically analyzed. The two solutions (before and
after adding hydrogen) were NMR spectroscopically identical. 1H NMR (400.17
MHz, CD3CN): δ = 1.11 (t, 9H, N(CH2CH3)3), 1.96 (quint., 1H, CHD2CN), 2.76 (q,
6H, N(CH2CH3)3), 4.74 (sept. 3H, B(Ohfip)3 adduct), 5.38 (sept. 3H, [XB(Ohfip)3]−.
19F NMR (376.54 MHz, CD3CN): δ = −76.0 (d, 18F, B(Ohfip)3 adduct), −75.4 (d,
18F, [XB(Ohfip)3]−). 11B NMR (376.54 MHz, CD3CN): δ = 3.2 (br, 1B,
[XB(Ohfip)3]−), 16.2 (br, 1B, B(Ohfip)3 adduct).
Analysis of the system 1/2,6-lutidine: Compound 1 (240 mg, 0.5 mmol, 1 eq.)
and 1,2-lutidine (50 mg, 0.5 mmol, 1 eq.) were mixed in a J. Young NMR tube
and dissolved in d3-acetonitrile. The reaction mixture was characterized by NMR
spectroscopy. Subsequently the reaction mixture was evacuated, exposed to
hydrogen pressure of one bar, and was then again NMR spectroscopically
analyzed. The two solutions (before and after adding hydrogen) were NMR
spectroscopically identical. 1H NMR (400.17 MHz, CD3CN): δ = 1.96 (quint., 1H,
51
CHD2CN), 2.47 (s, 6H, N(C6H3(2,6-CH3))), 5.47 (sept. 3H, B(Ohfip)3), 7.00 and
7.52 (m, 3H, N(C6H3(2,6-CH3))). 19F NMR (376.54 MHz, CD3CN): δ = −76.0 (d,
18F, B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN): δ = 17.4 (br, 1B, B(Ohfip)3).
Analysis of the system 1/NHPh2: Compound 1 (150 mg, 0.3 mmol, 1 eq.) and
NHPh2 (50 mg, 0.3 mmol, 1 eq.) were mixed in a J. Young NMR tube and
dissolved in d3-acetonitrile. The reaction mixture was characterized by NMR
spectroscopy. Subsequently the reaction mixture was evacuated, exposed to
hydrogen pressure of one bar, and was then again NMR spectroscopically
analyzed. The two solutions (before and after adding hydrogen) were NMR
spectroscopically identical. 1H NMR (400.17 MHz, CD3CN): δ = 1.96 (quint., 1H,
CHD2CN), 5.50 (sept. 3H, B(Ohfip)3), 6.80 (s, 1H, NHPh2), 6.94 (m, 2H,
NH(p)Ph2), 7.14 (m, 4H, NH(m)Ph2), 7.23 (m, 4H, NH(o)Ph2). 19F NMR (376.54
MHz, CD3CN): δ = −75.9 (d, 18F, B(Ohfip)3). 11B NMR (376.54 MHz, CD3CN): δ =
17.4 (br, 1B, B(Ohfip)3).
Analysis of the system 1/1,3-dimethyl-4,5-diphenyl imidazole-2-ylidene: Compound 1 (110 mg, 0.2 mmol, 1 eq.) and 1,3-dimethyl-4,5-diphenyl imidazole-
2-ylidene (45 mg, 0.2 mmol, 1 eq.) were mixed in a J. Young NMR tube and
dissolved in d8-toluene. Subsequently the yellow reaction mixture was evacuated,
exposed to hydrogen pressure of one bar, and NMR spectroscopically analyzed. 1H NMR (400.17 MHz, d8-toluene): δ = 2.09 (m, 1H, C6D5CHD2), 3.45 (s, 6H,
2xCH3 of the NHC compound) 5.12 (sept. 3H, B(Ohfip)3 adduct), 6.70 (m, 6H,
m/p protons of the NHC compound), 6.90 (m, 4H, o protons of the NHC
compound), 7.0 - 7.1 (m, aromatic proton of d8-toluene). 19F NMR (376.54 MHz,
d8-toluene): δ = −73.6 (d, 18F, B(Ohfip)3 adduct).11B NMR (376.54 MHz, CD3CN):
δ = 1.5 (br, 1B, B(Ohfip)3 adduct).
Analysis of the system 1/1,3-bis(2,6-diisopropylphenyl)imidazol-2-yliden: Compound 1 (60 mg, 0.1 mmol, 1 eq.) and 1,3-bis(2,6-
diisopropylphenyl)imidazol-2-yliden (45 mg, 0.1 mmol, 1 eq.) were mixed in a J.
Young NMR tube and dissolved in toluene. The brown reaction mixture was
characterized by NMR spectroscopy. Subsequently the reaction mixture was
evacuated, exposed to hydrogen pressure of one bar, and was then again NMR
spectroscopically analyzed. The two solutions (before and after adding hydrogen)
52
53
were NMR spectroscopically identical. 11B NMR (376.54 MHz, no lock in toluene):
δ = 1.9 (br, 1B, B(Ohfip)3 adduct).
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57
58
3.6 Appendix
Legend to the following tables and calculations herein.
BP86 BP86 density functional
h Hartree, 1H = 2625.5 kJ mol−1
Evrt Sum of vibrational, rotational and translational
energy including ZPE (= FreeH energy)
ZPE Zero point vibrational energy
S Entropy
H Enthalpy
G Gibbs Energy
G3 Energy G3 Energy including FreeH energy
CCSD(T) Coupled-cluster with single and double and
perturbative triple excitations
MP2 Møller-Plesset second-order perturbation
energy
Tot. En. OC corr. Total energy outlying charge corrected
For H+ and Cu+ the Sackur-Tetrode equation was used to calculate the entropy.
Table 3.7: Energetic data of the FLP screening. The thermal energy corrections are calculated at 298.15 K. NHC = 1,3-dimethyl-4,5-diphenyl imidazole-2-ylidene, IDipp = 1,3-bis(2,6-diisopropylphenyl)imidazol-2-yliden. Permittivity values are given in parenthesis.
Compound (RI-)
BP86/SV(P)
energy [h]
Evrt
[kJ mol−1]
Entropy
S
[kJ mol−1]
Enthalpy H
[kJ mol−1
Gibbs Energy
G
[kJ mol−1]
Tot. En. OC
corr. (8.93)
[h]
ΔSolvG0
(8.93)
[kJ mol−1]
Tot. En. OC
corr. (14.3)
[h]
ΔSolvG0
(14.3)
[kJ mol−1]
Tot. En. OC
corr. (36.64)
[h]
ΔSolvG0
(36.64)
[kJ mol−1]
1 −2390.95955 496.66 0.87631 −6276989.1 −6277250.4 −2390.96505 −6.5 −2390.96545 −7.5 −2390.96589 −8.7
B(C6F5)3 −2206.75293 471.63 0.82219 −5793377.8 −5793622.9 −2206.75759 −4.3 −2206.75793 −5.2 −2206.75829 −6.1
H2 −1.16996 31.70 0.13689 −3037.6 −3078.4 −1.17015 7.4 −1.17017 7.4 −1.17018 7.4
[HB(Ohfip)3]− −2391.59560 513.08 0.85226 −6278642.6 −6278896.7 −2391.65060 −136.4 −2391.65449 −146.6 −2391.65861 −157.5
[HB(C6F5)3] − −2207.45157 492.68 0.83844 −5795191.0 −5795441.0 −2207.49784 −113.5 −2207.50108 −122.0 −2207.50452 −131.1
PPh3 −1035.73393 741.32 0.56281 −2718586.0 −2718753.8 −1035.74220 −13.7 −1035.74286 −15.5 −1035.74357 −17.3
[HPPh3]+ −1036.11240 769.91 0.57343 −2719551.1 −2719722.0 −1036.16659 −134.3 −1036.17041 −144.4 −1036.17447 −155.0
PtBu3 −814.40529 983.86 0.53982 −2137242.9 −2137403.8 −814.40883 −1.3 −814.40912 −2.1 −814.40944 −2.9
[HPtBu3]+ −814.80160 1013.94 0.54805 −2138253.3 −2138416.7 −814.85664 −136.5 −814.86045 −146.5 −814.86447 −157.1
NEt3 −292.17060 547.89 0.39098 −766546.5 −766663.0 −292.17248 3.0 −292.17262 2.7 −292.17277 2.3
[HNEt3]+ −292.55037 588.17 0.39266 −767503.3 −767620.4 −292.61891 −172.0 −292.62368 −184.5 −292.62873 −197.8
2,6-lutidine −326.68165 386.07 0.37027 −857317.4 −857427.8 −326.68679 −5.5 −326.68720 −6.6 −326.68765 −7.8
[2,6-lutidineH] −327.06021 422.08 0.37479 −858275.3 −858387.0 −327.12863 −171.7 −327.13338 −184.2 −327.13841 −197.4
NHPh2 −518.28400 533.52 0.42554 −1360223.8 −1360350.7 −518.29259 −14.6 −518.29326 −16.4 −518.29399 −18.3
[HNHPh2]+ −518.63119 569.39 0.43935 −1361099.5 −1361230.5 −518.70241 −179.0 −518.70757 −192.6 −518.71307 −207.0
NHC −766.36360 780.61 0.54762 −2011312.2 −2011475.5 −766.37643 −25.7 −766.37747 −28.5 −766.37861 −31.4
[HNHC] + −766.80363 816.09 0.55535 −2012432.0 −2012597.6 −766.86254 −146.7 −766.86680 −157.9 −766.87135 −169.8
IDipp −1159.13877 1532.87 0.79523 −3041795.1 −3042032.2 −1159.14887 −18.6 −1159.14973 −20.8 −1159.15066 −23.3
[HIDipp]+ −1159.57973 1566.16 0.79153 −3042919.5 −3043155.5 −1159.62947 −122.6 −1159.63299 −131.9 −1159.63671 −141.7
59
Table 3.8: Part 1 of the energetic data of the validation study. The thermal energy corrections are calculated at 298.15 K.
Compound CCSD(T)/AVDZ energy
[h]
MP2/AVDZ energy
[h]
MP2/AVQZ energy
[h]
BP86/def-TZVP energy
[h]
Evrt (BP86/def-TZVP)
[kJ mol−1
Cl− −459,74380 −459.72656 −459.80308 −460.32339 3.72
H− −0.52403 −0.51196 −0.51709 −0.51016 3.72
F− −99.66863 −99.66595 −99.77441 −99.87911 3.72
BF3 −323.92630 −323.90489 −324.27979 −324.70276 40.86
[BF3Cl]− −783.73105 −783.69171 −784.13919 −785.08739 45.51
[BF3H]− −324.57583 −324.54755 −324.92064 −325.35264 61.19
[BF3F]− −423.72669 −423.69914 −424.18275 −424.72498 47.76
BCl3 −1403.92243 −1403.86501 −1404.12135 −1405.75626 31.04
[BCl3Cl]− −1863.73645 −1863.66506 −1864.00194 −1866.15534 37.12
[BCl3H]− −1404.60198 −1404.53974 −1404.80295 −1406.44347 54.70
[BCl3F]− −1503.73570 −1503.67417 −1504.04666 −1505.80099 53.50
BBr3 −1271.76505 −1271.71898 −1271.95235 −7748.28799 28.48
[BBr3Cl]− −1731.58766 −1731.52786 −1731.84258 −8208.69430 34.98
[BBr3H]− −1272.45777 −1272.40631 −1272.65045 −7748.98759 52.47
[BBr3F]− −1371.58573 −1371.53548 −1371.88830 −7848.34143 38.06
AlF3 −541.21627 −541.19893 −541.55605 −542.33712 31.10
[AlF3Cl]− −1001.07365 −1001.03787 −1001.47532 −1002.78145 37.90
[AlF3H]− −541.89719 −541.87139 −542.22845 −543.02669 47.83
[AlF3F]− −641.06636 −641.04300 −641.51147 −642.40837 39.10
AlCl3 −1621.22805 −1621.17230 −1621.43538 −1623.44505 26.33
[AlCl3Cl]− −2081.08910 −2081.01786 −2081.36279 −2083.89330 33.90
[AlCl3H]− −1621.91982 −1621.85769 −1622.12563 −1624.14546 44.29
[AlCl3F]− −1721.08473 −1721.02497 −1721.40059 −1723.52361 35.23
60
Table 3.9: Part 2 of the energetic data of the validation study. The thermal energy corrections are calculated at 298.15 K.
Compound CCSD(T)/AVDZ energy
[h]
MP2/AVDZ energy
[h]
MP2/AVQZ energy
[h]
BP86/def-TZVP energy
[h]
Evrt (BP86/def-TZVP)
[kJ mol−1
AlBr3 −1489.08371 −1489.03864 −1489.29605 −7965.99036 24.82
[AlBr3Cl]− −1948.94503 −1948.88480 −1949.22499 −8426.43807 32.48
[AlBr3H]− −1489.77765 −1489.72620 −1489.98537 −7966.69272 43.01
[AlBr3F]− −1588.94082 −1588.89187 −1589.26271 −8066.06903 33.88
GaF3 −558.44397 −558.47027 −558.78456 −2224.91244 28.44
[GaF3Cl]− −1018.31462 −1018.32170 −1018.70754 −2225.62007 35.47
[GaF3H]− −559.14399 −559.16445 −559.48106 −2324.97033 45.77
[GaF3F]− −658.28543 −658.30518 −658.72602 −2324.97033 36.47
GaCl3 −1638.53793 −1638.52675 −1638.73494 −3306.07616 24.87
[GaCl3Cl]− −2098.39939 −2098.37088 −2098.65226 −3766.51500 32.36
[GaCl3H]− −1639.23295 −1639.21693 −1639.43356 −3306.78232 43.18
[GaCl3F]− −1738.37154 −1738.35446 −1738.67054 −3406.13003 33.41
GaBr3 −1506.00214 −1505.98457 −1506.23389 −9648.63444 23.68
[GaBrl3Cl]− −1965.86085 −1965.82640 −1966.15032 −10109.07026 31.31
[GaBr3H]− −1506.69587 −1506.67316 −1506.93225 −9649.33927 42.11
[GaBr3F]− −1605.83307 −1605.80988 −1606.16841 −9748.68578 32.36
PF5 −839.33115 −839.29605 −839.93790 −840.96036 55.75
[PF5Cl]− −1299.14595 −1299.09337 −1299.80322 −1301.35977 60.99
[PF5H]− −840.02963 −839.98853 −840.61981 −841.66526 79.23
[PF5F]− −939.15165 −939.11028 −939.85401 −941.00075 63.25
61
Table 3.10: Part 3 of the energetic data of the validation study. The thermal energy corrections are calculated at 298.15 K.
Compound CCSD(T)/AVDZ energy
[h]
MP2/AVDZ energy
[h]
MP2/AVQZ energy
[h]
BP86/def-TZVP energy
[h]
Evrt (BP86/def-TZVP)
[kJ mol−1
PCl5 −2639.30190 −2639.21138 −2639.67582 −2642.74701 42.67
[PCl5Cl]− −3099.11408 −3099.01083 −3099.55194 −3103.14468 49.53
[PCl5H]− −2640.01841 −2639.92389 −2640.38628 −2643.47134 66.97
[PCl5F]− −2739.11992 −2739.02599 −2739.60030 −2742.78836 51.66
AsF5 −829.99443 −829.96364 −830.51661 −2735.65774 49.49
[AsF5Cl]− −1289.82762 −1289.77642 −1290.40684 −3196.08450 55.69
[AsF5H]− −830.70791 −830.67175 −831.22162 −2736.38634 70.83
[AsF5F]− −929.83166 −929.79451 −930.45322 −2835.71345 57.30
B(OH)3 −251.94592 −251.91094 −252.20599 −252.59931 135.64
[B(OH)3H]− −252.53693 −252.49575 −252.79510 −253.19643 152.71
[B(OH)3F]− −351.68837 −351.64777 −352.05589 −352.56773 140.10
SbF5 −738.69619 −738.67462 −739.20825 −504.96124 46.33
[SbF5Cl]− −1198.57085 −1198.53137 −1199.13783 −965.41970 53.39
[SbF5H]− −739.42544 −739.39885 −739.93919 −505.71266 65.59
[SbF5F]− −838.55453 −838.52653 −839.16811 −605.04171 54.62
[SbF4]− −639.15362 −639.14524 −639.57548 −405.23423 34.55
HF −100.26360 −100.25575 −100.36976 −100.49251 29.93
62
Table 3.11: Energetic data of the reference systems. The thermal energy corrections are calculated at 298.15 K.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
G3 Enthalpy H
[h]
MeSi3+ −408.80757 297.89 0.36 −1073028.00 −1073136.05 −408.67945
MeSi3Cl −869.25828 309.46 0.36 −2281934.36 −2282042.36 −869.08993
MeSi3H −409.68913 322.45 0.33 −1075317.98 −1075416.84 −409.56127
MeSi3F −508.92757 311.07 0.35 −1335880.88 −1335985.90 −508.85133
MeSi3Me −448.98131 399.76 0.36 −1178402.69 −1178508.71 −448.84834
Cl− −460.12124
H− −0.51636
F− −99.80683
Me− −39.78805
63
Table 3.12: Part 1 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
BF3 −324.31769 41.36 0.26 −851455.49 −851531.69
[BF3Cl]− −784.53230 46.60 0.30 −2059748.31 −2059836.54
[BF3H]− −324.94632 62.26 0.27 −853085.06 −853164.20
[BF3F]− −424.20099 49.21 0.27 −1113692.26 −1113772.75
[BF3Me]− −364.24312 136.67 0.30 −956184.81 −956273.81
BCl3 −1405.27457 31.32 0.29 −3689528.65 −3689615.38
[BCl3Cl]− −1865.50372 37.75 0.32 −4897858.45 −4897954.10
[BCl3H]− −1405.93901 54.79 0.30 −3691249.66 −3691340.48
[BCl3F]− −1505.18222 40.74 0.32 −3951827.76 −3951923.42
[BCl3Me]− −1445.23236 129.98 0.33 −3794339.55 −3794439.29
BBr3 −7747.33631 28.68 0.33 −20340677.81 −20340774.90
[BBr3Cl]− −8207.57694 35.40 0.37 −21549037.46 −21549147.68
[BBr3H]− −7748.01880 52.50 0.34 −20342445.85 −20342547.76
[BBr3F]− −8064.92779 34.35 0.39 −21174511.73 −21174626.78
[BBr3Me]− −8004.96231 121.67 0.41 −21016984.43 −21017105.99
BI3 −59.27797 26.95 0.35 −155605.47 −155710.23
[BI3Cl]− −519.53720 34.45 0.40 −1364013.19 −1364131.35
[BI3H]− −59.98640 51.22 0.37 −157441.19 −157550.69
[BI3F]− −159.21958 37.55 0.39 −417992.56 −418107.67
[BI3Me]− −99.27605 127.33 0.40 −260520.46 −260638.71
AlF3 −541.88983 31.84 0.28 −1422702.84 −1422786.10
[AlF3Cl]− −1002.16640 39.20 0.32 −2631156.24 −2631251.75
64
Table 3.13: Part 2 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[AlF3H]− −542.56444 49.17 0.29 −1424456.72 −1424542.53
[AlF3F]− −641.82254 40.73 0.30 −1685068.30 −1685156.52
[AlF3Me]− −581.85677 126.79 0.33 −1527541.49 −1527640.78
AlCl3 −1622.90465 26.53 0.31 −4260923.39 −4261017.25
[AlCl3Cl]− −2083.18568 34.24 0.35 −5469388.11 −5469491.74
[AlCl3H]− −1623.58970 44.67 0.33 −4262703.86 −4262801.44
[AlCl3F]− −1722.84767 35.87 0.35 −4523315.44 −4523418.76
[AlCl3Me]− −1662.88217 122.80 0.37 −4365788.48 −4365898.83
AlBr3 −7964.98011 24.94 0.35 −20912107.52 −20912211.62
[AlBr3Cl]− −8425.26423 32.77 0.41 −22120580.25 −22120701.18
[AlBr3H]− −7965.67055 43.18 0.37 −20913902.03 −20914011.23
[AlBr3F]− −8004.96231 34.35 0.39 −21174511.73 −21174626.78
[AlBr3Me]− −8064.92779 121.67 0.41 −21016984.43 −21017105.99
AlI3 −276.93301 24.13 0.38 −727063.77 −727175.64
[AlI3Cl]− −737.22532 32.15 0.42 −1935557.82 −1935683.90
[AlI3H]− −277.63619 42.44 0.39 −728891.68 −729008.42
[AlI3F]− −376.89045 33.87 0.41 −989493.31 −989615.57
[AlI3Me]− −316.92785 121.31 0.43 −831973.45 −832102.32
GaF3 −2224.30029 29.01 0.29 −5839891.16 −5839978.55
[GaF3Cl]− −2684.57647 36.21 0.34 −7048343.68 −7048444.36
[GaF3H]− −2224.98949 46.42 0.30 −5841683.27 −5841773.85
[GaF3F]− −2324.21882 37.40 0.31 −6102219.87 −6102313.54
65
Table 3.14: Part 3 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[GaF3Me]− −2264.27778 125.00 0.35 −5944756.47 −5944861.69
GaCl3 −3305.36839 24.94 0.33 −8678250.35 −8678348.11
[GaCl3Cl]− −3765.64032 32.51 0.36 −9886691.33 −9886799.41
[GaCl3H]− −3306.05889 43.35 0.34 −8680044.84 −8680146.49
[GaCl3F]− −3405.28715 33.76 0.36 −8940579.23 −8940687.11
[GaF3Me]− −3345.34742 122.10 0.39 −8783118.53 −8783235.71
GaBr3 −9647.45623 23.70 0.36 −25329466.62 −25329574.65
[GaBr3Cl]− −10107.72841 31.36 0.41 −26537908.17 −26538031.29
[GaBr3H]− −9648.14898 42.11 0.38 −25331267.04 −25331380.42
[GaBr3F]− −9747.37651 32.60 0.40 −25591799.42 −25591918.95
[GaBr3Me]− −9687.43727 121.25 0.43 −25434339.69 −25434469.31
GaI3 −1959.42418 23.20 0.39 −5144462.09 −5144577.66
[GaI3Cl]− −2419.70219 31.05 0.44 −6352918.77 −6353049.13
[GaI3H]− −1960.12637 41.65 0.40 −5146287.25 −5146407.92
[GaI3F]− −2059.35144 32.29 0.43 −5406813.02 −5406939.94
[GaI3Me]− −1999.41466 121.12 0.45 −5249359.59 −5249492.62
PF5 −840.24920 57.03 0.31 −2206023.18 −2206115.23
[PF5Cl]− −1300.47676 63.16 0.33 −3414349.10 −3414448.92
[PF5H]− −840.92368 80.76 0.31 −2207770.29 −2207862.21
[PF5F]− −940.15393 65.88 0.30 −2468315.19 −2468405.78
[PF5Me]− −880.21459 155.35 0.36 −2310854.38 −2310961.56
PCl5 −2641.86516 43.04 0.38 −6936197.89 −6936309.79
66
Table 3.15: Part 4 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[PCl 5Cl]− −3102.09251 49.86 0.42 −8144522.57 −8144648.16
[PCl 5H]− −2642.56490 66.94 0.40 −6938011.16 −6938129.56
[PCl 5F]− −2741.76810 52.36 0.41 −7198484.71 −7198606.02
[PCl 5Me]− −2681.85046 143.36 0.44 −7041079.36 −7041210.43
AsF5 −2734.77740 50.42 0.33 −7180132.50 −7180230.72
[AsF5Cl]− −3195.03263 57.06 0.36 −8388530.58 −8388637.57
[AsF5H]− −2735.47713 72.02 0.33 −7181948.05 −7182046.65
[AsF5F]− −2834.69432 58.86 0.33 −7442456.95 −7442555.10
[AsF5Me]− −2774.76975 148.93 0.39 −7285034.32 −7285149.74
SbF5 −504.34073 47.06 0.35 −1324102.08 −1324207.90
[SbF5Cl]− −964.63043 54.24 0.39 −2532590.12 −2532705.74
[SbF5H]− −505.06879 66.27 0.36 −1325994.40 −1326100.42
[SbF5F]− −604.28148 55.55 0.36 −1586489.03 −1586596.20
[SbF5Me]− −544.36348 145.04 0.42 −1429084.23 −1429209.53
[SbF4]− −404.71210 35.45 0.35 −1062537.73 −1062642.09
HF −100.34485 28.63 0.17 −263425.30 −263477.25
MeF −139.62734 107.38 0.22 −366483.12 −366549.63
B(OH)3 −252.27649 133.32 0.28 −662218.64 −662475.47
[B(OH)3H]− −252.85454 150.85 0.29 −663718.80 −663718.80
[B(OH)3F]− −352.10612 138.48 0.31 −924317.18 −924317.18
[B(OH)3Me]− −292.14938 225.69 0.32 −766812.95 −766897.76
B(Ohfip)3 −2390.95983 496.86 0.88 −6276989.59 −6277251.89
67
Table 3.16: Part 5 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[B(Ohfip)3Cl]− −2851.17065 496.54 0.87 −7485278.03 −7485537.36
[B(Ohfip)3H]− −2391.60549 513.18 0.87 −6278668.47 −6278927.43
[B(Ohfip)3F]− −2490.85720 500.15 0.89 −6539267.86 −6539531.90
[B(Ohfip)3Me]− −2430.89606 587.97 0.88 −6381751.46 −6382013.04
Al(OC(CF3)3)3 −3618.91753 540.66 1.06 −9500961.01 −9501277.13
[Al(OC(CF3)3)3Cl]− −4079.21138 548.19 1.13 −10709459.60 −10709797.32
[Al(OC(CF3)3)3H]− −3619.61749 558.40 1.11 −9502781.03 −9503112.16
[Al(OC(CF3)3)3F]− −3718.87763 549.80 1.12 −9763398.12 −9763732.36
[Al(OC(CF3)3)3Me]− −3658.91005 636.25 1.14 −9605866.20 −9606207.05
B(OTeF5)3 −1771.19399 206.58 0.78 −4650078.46 −4650312.34
[B(OTeF5)3Cl]− −2231.47676 211.90 0.85 −5858550.17 −5858803.31
[B(OTeF5)3H]− −1771.92090 228.68 0.83 −4651964.89 −4652213.21
[B(OTeF5)3F]− −1871.15773 215.61 0.83 −4912525.24 −4912772.12
[B(OTeF5)3Me]− −1811.21457 304.14 0.85 −4755055.33 −4755307.57
As(OTeF5)3 −5146.23685 329.24 1.24 −13511164.60 −13511534.16
[As(OTeF5)3Cl]− −5606.54953 335.27 1.24 −14719714.10 −14720083.44
[As(OTeF5)3H]− −5147.02237 350.71 1.22 −13513205.50 −13513570.11
[As(OTeF5)3F]− −5246.21827 337.96 1.22 −13773658.10 −13774022.65
[As(OTeF5)3Me]− −5186.31577 428.80 1.25 −13616292.65 −13616665.75
Sb(OTeF5)3 −2915.82616 326.64 1.27 −7655201.63 −7655580.10
[Sb(OTeF5)3Cl]− −3376.16324 334.10 1.29 −8863813.77 −8864197.49
[Sb(OTeF5)3H]− −2916.62459 346.36 1.26 −7657278.19 −7657654.19
68
Table 3.17: Part 6 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[Sb(OTeF5)3F]− −3015.81754 335.75 1.29 −7917720.88 −7918106.00
[Sb(OTeF5)3Me]− −2955.92663 426.83 1.31 −7760385.61 −7760777.44
B(CF3)3 −1037.00986 152.98 0.50 −2722524.30 −2722672.16
[B(CF3)3Cl]− −1497.30582 160.10 0.50 −3931028.83 −3931178.10
[B(CF3)3H]− −1037.74761 176.43 0.46 −2724437.81 −2724575.21
[B(CF3)3F]− −1136.97516 162.34 0.48 −2984974.84 −2985119.13
[B(CF3)3Me]− −1077.03536 251.85 0.49 −2827512.78 −2827659.65
B(C6F5)3 −2206.75271 471.95 0.81 −5793376.89 −5793618.43
[B(C6F5)3Cl]− −2667.00121 476.04 0.85 −7001759.82 −7002011.93
[B(C6F5)3H]− −2207.45155 492.67 0.84 −5795190.96 −5795442.86
[B(C6F5)3F]− −2306.67696 477.88 0.85 −6055723.05 −6055975.83
[B(C6F5)3Me]− −2246.72740 568.12 0.85 −5898234.66 −5898489.38
Al(C6F5)3 −2424.31626 462.17 0.88 −6364601.94 −6364863.97
[Al(C6F5)3Cl]− −2884.60755 466.90 0.91 −7573096.59 −7573368.66
[Al(C6F5)3H]− −2425.01252 477.22 0.89 −6366414.92 −6366679.22
[Al(C6F5)3F]− −2524.27281 468.18 0.90 −6627032.84 −6627300.52
[Al(C6F5)3Me]− −2464.30313 555.74 0.93 −6469494.28 −6469772.08
Ga(C6F5)3 −4106.80716 461.49 0.89 −10781999.30 −10782265.18
[Ga(C6F5)3Cl]− −4567.08293 466.51 0.91 −11990452.92 −11990724.71
[Ga(C6F5)3H]− −4107.50218 476.60 0.89 −10783808.97 −10784075.12
[Ga(C6F5)3F]− −4206.73211 467.52 0.91 −11044347.22 −11044617.21
[Ga(C6F5)3Me]− −4146.78826 555.62 0.92 −10886875.95 −10887150.01
69
Table 3.18: Part 7 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
B(C12F9)3 −4089.39616 909.52 1.34 −10735838.51 −10736237.10
[B(C12F9)3Cl]− −4549.62632 911.86 1.38 −11944175.06 −11944585.04
[B(C12F9)3H]− −4090.08244 929.20 1.34 −10737620.66 −10738021.05
[B(C12F9)3F]− −4189.31221 914.55 1.37 −10998164.07 −10998571.88
[B(C12F9)3Me]− −4129.34508 1005.86 1.36 −10840628.46 −10841033.83
B(C6H3-3,5-(CF3)2)3 −2740.16520 879.48 1.05 −7193449.18 −7193762.90
[B(C6H3-3,5-(CF3)2)3Cl]− −3200.43072 883.35 1.09 −8401877.03 −8402203.18
[B(C6H3-3,5-(CF3)2)3H]− −2740.86442 898.74 1.06 −7195265.73 −7195582.83
[B(C6H3-3,5-(CF3)2)3F]− −2840.10065 884.94 1.09 −7455825.25 −7456150.07
[B(C6H3-3,5-(CF3)2)3Me]− −2780.14741 974.12 1.09 −7298328.22 −7298652.37
B(C10F7)3 −3262.34797 741.74 1.10 −8564583.00 −8564909.79
[B(C10F7)3Cl]− −3722.60767 745.92 1.12 −9772995.25 −9773329.56
[B(C10F7)3H]− −3263.06000 762.19 1.11 −8566432.00 −8566764.12
[B(C10F7)3F]− −3362.28421 747.82 1.12 −8826960.51 −8827294.29
[B(C10F7)3Me]− −3302.33656 838.17 1.12 −8669477.01 −8669810.85
F4C6(1,2-(B(C6F5)2)2 −3586.44641 784.90 1.20 −9415463.52 −9415820.57
[F4C6(1,2-(B(C6F5)2)2Cl]− −4046.72254 789.20 1.23 −10623918.81 −10624286.18
[F4C6(1,2-(B(C6F5)2)2H]− −3587.16490 805.00 1.21 −9417329.83 −9417689.56
[F4C6(1,2-(B(C6F5)2)2F]− −3686.39808 791.83 1.23 −9677880.70 −9678247.15
[F4C6(1,2-(B(C6F5)2)2Me]− −3626.42419 879.05 1.22 −9520331.44 −9520695.13
B2(C6F5)2(C6F4)2 −2759.39574 620.04 0.99 −7244198.58 −7244494.68
[B2(C6F5)2(C6F4)2Cl]− −3219.65734 623.98 1.02 −8452616.07 −8452920.01
70
Table 3.19: Part 8 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[B2(C6F5)2(C6F4)2H]− −2760.10557 639.65 1.00 −7246042.64 −7246339.39
[B2(C6F5)2(C6F4)2F]− −2859.32945 625.80 1.01 −7506569.78 −7506872.16
[B2(C6F5)2(C6F4)2Me]− −2799.38597 716.35 1.01 −7349097.03 −7349397.63
PF2(C2F5)3 −2265.88419 287.78 0.70 −5948811.34 −5949021.39
[PF2(C2F5)3Cl]− −2726.10290 292.57 0.69 −7157115.37 −7157322.52
[PF2(C2F5)3H]− −2266.56226 309.87 0.68 −5950569.54 −5950773.36
[PF2(C2F5)3F]− −2365.78476 294.99 0.69 −6211094.08 −6211299.02
[PF2(C2F5)3Me]− −2305.83768 385.36 0.71 −6053612.04 −6053823.61
B(CN)3 −303.16014 92.44 0.35 −795855.06 −795959.01
[B(CN)3Cl]− −763.45377 100.38 0.38 −2004352.65 −2004465.61
[B(CN)3H]− −303.89945 116.56 0.35 −797772.02 −797875.46
[B(CN)3F]− −403.12354 102.33 0.37 −1058300.09 −1058410.04
[B(CN)3Me]− −343.18480 192.85 0.38 −900839.80 −900953.04
B11Cl11CH2TMS −5781.55702 534.11 0.81 −15178999.18 −15179241.68
[B11Cl11CH2TMSCl]− −6241.86871 543.36 0.85 −16387542.89 −16387796.79
[B11Cl11CH2TMSH]− −5782.29716 556.44 0.82 −15180920.08 −15181165.16
[B11Cl11CH2TMSF]− −5881.53834 544.90 0.84 −15441490.36 −15441744.27
[B11Cl11CH2TMSMe]− −5821.58782 633.59 0.86 −15284000.97 −15284257.79
TMS-F-Al(OC(CF3)3)3 −4127.88010 861.40 1.28 −10836926.59 −10837308.52
[FAl(OC(CF3)3)3]− −3718.87761 549.78 1.13 −9763398.10 −9763733.57
TMSCl −869.25828 309.46 0.36 −2281934.36 −2282042.36
TMSH −409.68913 322.45 0.33 −1075317.98 −1075416.84
71
Table 3.20: Part 9 of the energetic data to determine the affinity values. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
TMSF −508.92757 311.07 0.35 −1335880.88 −1335985.90
TMSMe −448.98131 399.76 0.36 −1178402.69 −1178508.71
[ClAl(OC(CF3)3)3]− −4079.21128 547.89 1.14 −10709459.64 −10709798.23
[HAl(OC(CF3)3)3]− −3619.61838 558.38 1.11 −9502783.39 −9503114.66
[MeAl(OC(CF3)3)3]− −3658.91052 636.28 1.15 −9605867.39 −9606210.06
Table 3.21: Part 1 of the energetic data to determine the decomposition values of WCAs. The thermal energy corrections are calculated at 298.15 K. The energetic data of [BX4]−, [AlX4]−, GaX4]−, [AsF6]− and [SbF6] − (X = F, Cl, Br, I) and all underlying Lewis acids were already shown in the tables above.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
H+ 0 6.20 0.11 8.68 −23.78
Cu+ −1640.21528 6.20 0.16 −4306392.95 −4306440.82
[B(Ohfip)4]− −3179.74703 650.70 1.07 −8347804.45 −8348124.72
[Al(OC(CF3)3)4]− −4744.54973 719.65 1.37 −12456140.63 −12456548.10
[B(OTeF5)4]− −2353.44290 270.92 0.97 −6178714.47 −6179002.48
[As(OTeF5)6]− −5728.50557 393.29 1.40 −15039852.88 −15040269.02
[Sb(OTeF5)6]− −3498.11286 391.67 1.45 −9183936.14 −9184368.77
[B(CN)4]− −396.10083 121.24 0.37 −1039842.97 −1039953.51
[B(CF3)4]− −1374.57309 200.53 0.54 −3608752.38 −3608913.48
72
Table 3.22: Part 2 of the energetic data to determine the decomposition values of WCAs. The thermal energy corrections are calculated at 298.15 K.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
[B(C6F5)4]− −2934.16744 622.30 0.99 −7703061.17 −7703356.33
[Al(C6F5)4]− −3151.74555 612.02 1.06 −8274324.95 −8274641.47
[Ga(C6F5)4]− −4834.22867 611.18 1.07 −12691702.06 −12692020.36
[B(C12F9)4]− −5444.37581 1206.18 1.69 −14293054.48 −14293558.79
[B(C6H3(CF3)2)4]− −3645.38949 1164.52 1.33 −9569839.56 −9570236.12
[B(C10F7)4]− −4341.64150 981.78 1.36 −11398038.93 −11398444.87
[F4C6 (1,2-
(B(C6F5)2)2)(C6F5)]− −4313.83839 934.56 1.37 −11325088.80 −11325497.31
[B2(C6F5)3(C6F4)2]− −3486.81296 770.04 1.16 −9153889.78 −9154236.66
HF −100.34485 28.62 0.17 −263425.31 −263477.26
HCl −460.69877 22.83 0.19 −1209543.93 −1209599.79
HBr −2574.72899 21.24 0.20 −6759953.00 −6760012.37
HI −12.05201 19.44 0.21 −31620.75 −31682.48
HOhfip −789.24620 183.27 0.40 −2071988.04 −2072108.29
HOC(CF3)3 −1126.04299 202.27 0.47 −2956232.37 −2956372.04
HOTeF5 −582.66184 86.41 0.39 −1529695.59 −1529811.43
HCN −93.35344 48.13 0.20 −245049.78 −245110.04
HCF3 −337.99423 74.06 0.26 −887330.69 −887408.36
HC6F5 −727.89270 176.03 0.38 −1910911.04 −1911025.17
HC12F9 −1355.44401 322.06 0.57 −3558407.25 −3558576.64
HC6(H)3(CF3)2 −905.70077 312.23 0.46 −2377611.72 −2377747.49
HC10F7 −1079.75991 266.17 0.48 −2834651.80 −2834795.35
73
74
Table 3.23: Part 3 of the energetic data to determine the decomposition values of WCAs. The thermal energy corrections are calculated at 298.15 K.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G[kJ mol−1]
CuF −1740.29714 10.31 0.23 −4569154.76 −4569222.28
CuCl −2100.67728 9.45 0.24 −5515337.28 −5515408.07
CuBr −4214.71358 9.08 0.25 −11065761.08 −11065835.30
CuI −1652.04406 8.98 0.26 −4337446.73 −4337523.14
CuOhfip −2429.19188 158.60 0.45 −6377706.49 −6377839.64
CuOC(CF3)3 −2765.99722 178.06 0.51 −7261972.82 −7262125.20
CuOTeF5 −2222.64055 64.00 0.43 −5835498.52 −5835627.42
CuCN −1733.31229 28.33 0.25 −4550797.93 −4550872.82
CuCF3 −1977.92109 46.83 0.31 −5193002.29 −5193093.45
CuC6F5 −2367.83452 152.82 0.43 −6216617.91 −6216746.21
CuC12F9 −2995.38543 298.74 0.62 −7864113.17 −7864296.66
CuC6H3(CF3)2 −2545.63221 288.37 0.50 −6683291.97 −6683440.99
CuC10F7 −2719.70140 242.77 0.53 −7140357.96 −7140515.62
4 Synthesis and Oxidation of Hexafluoroisopropoxyborylferrocenes
A part of this chapter was published in Pure and Applied Chemistry and was
reproduced:
http://dx.doi.org/10.1351/PAC-CON-12-02-14
Numbering and nomenclature of subchapters, figures, tables and compounds
were changed to be consistent in this thesis. Some information was added for
clarity and completeness. The technical data for all executed calculations is
deposited on the file server of the group of Prof. Krossing (path:
public\Ehemalige\Hannes Böhrer 2014\SI_Disseratation\SI_Tech_Chapt4.pdf).
4.1 Introduction
Ferrocenylboranes were the subject of recent research and found widespread
application in chemistry. The main advantage of the ferrocene backbone can be
assigned to its electrochemical properties paired with the chemical and thermal
inertness. Thus, ferrocenylboranes function, for example, as anion sensors for
selective fluoride or hydrogen fluoride detection.[1] Owning to the relatively weak
Lewis acidity of dialkoxyborylferrocenes, their fluoride binding is weak and
recognition by NMR or chemical monitoring is possible even in the presence of
other potentially competing anions. The ferrocenylborates on the basis of
tris(1-pyrazolyl)borates can be assembled to heterooligometallic complexes and
polymers to explore novel magnetic and conducting compounds.[2]
Ferrocenylboranes linked via B-N bonds and bifunctional pyridine bases
(4,4’-bipyridene or pyrazine) lead to charge-transfer polymers.[3] The charge-
transfer interaction occurs between the ferrocene donor and the electron-poor B-
N adduct bridge. In 1,3-dibora[3]ferrocenophanes, the two Cp rings of ferrocene
are connected by a B-E-B (E = O, S, Se, Te or NMe) fragment.[4] In these
compounds, the Cp rings bent to the heteroelements and are exactly staggered,
which is unusual for ferrocenophanes. The spontaneous formation of B-N bonds
was amongst others exploited in electron sponges,[5] which are composed of
75
ferrocene and up to four 2,2’-bipyridylboronium moieties. Those stable multi-step
redox systems are able to store up to nine electrons and the ligands are
electronically communicating. Oligonuclear metallocene aggregates like
[1.1]-diborataferrocenylphosphanes are capable of trapping Li ions.[6] In this
redox-switchable Li scavenger, the ion is not coordinating by main-group Lewis
bases. Bifunctional Lewis acids like Carpenters 1,1’-[B(C6F5)2]2Fc are highly
electrophilic, also 1,2-bifunctional ferrocenylboranes-based Lewis acids are
known.[7] Ansa-ferrocenes take advantage of the self-assembly and reversible
breaking of B-N or B-P Lewis acid-base pairs as interannular bridge. Thereby, the
structure is switchable between a rigid ansa-structure and a flexible open-chain
conformation by temperature or substitution pattern.[8]
One of the reasons for this abundance of different ferrocene-based boron
compounds is the straightforward preparation of dibromoborylferrocene and
1,1’-bis(dibromoboryl)ferrocene,[9] which are the basic synthesis modules for
many ferrocenylboranes. Another intriguing aspect of ferrocenylboranes based
chemistry is the still not well understood through-space B-Fe interaction, which
was part of detailed investigations.[10, 11] This interaction is linked to the bending
of the boryl residues toward the Fe center: the α* dip angle (α* = 180° − α, with α
being the angle Cp(centroid)-Cipso-B) is used as an indicator for the Lewis acidity
of the boryl moiety.
α*CipsoCOG
B
Fe
Figure 4.1: Definition of the dip angle α* and further geometric terms.
76
The higher the borylation degree of the ferrocenylboranes, the smaller α*. The
introduction of π-donating residues or adduct formation effect the same.
Ferrocenylboranes are redox-active and can be oxidized to the respective
ferrocinium compounds that should exert higher Lewis acidity. The first structural
analysis of a redox-switchable system was published recently (2008) for
9-ferrocenyl-9-borafluorene.[12] Interestingly, the higher Lewis acidity of this
compound is here accompanied by smaller values of the dip angle α*.
In this study we synthesized and characterized novel ferrocenylboranes and also
studied their subsequent oxidation. To combine a decent stability with sufficient
Lewis acidity, we chose the fluorinated sequent oxidation. To combine a decent
stability with sufficient Lewis acidity, we chose the fluorinated alkoxide
1,1,1,3,3,3-hexafluoro-2-propoxy ligand (Ohfip) for the boryl residue. The Ohfip
ligand was proven to provide suitable Lewis acidity to an electron-deficient central
atom, e.g., in the related B(Ohfip)3, as well as in monomeric Al[OC(CF3)3]3 or
dimeric Al2(Ohfip)6.[13, 14] Carpenter and coworkers suggested that oxidation
enhances the Lewis acidity of ferrocenylboranes, so that the B center in certain
borylferrocinium compounds abstracts a fluoride from [BF4]− counteranion.[15] To
avoid this side reaction, we decided to use the silver salt of the perfluorinated
alkoxyaluminate anion [Al(ORF)4]− (RF = C(CF3)3) as an oxidant.[14] The CF bonds
of the latter are more resistant against fluoride abstraction. Owing to its
delocalized negative charge, the [Al(ORF)4]− anion belongs to the category of
weakly coordinating anions (WCAs), which are ideal to investigate the crystal
structures of the oxidized compounds due to the minimized interaction between
the oppositely charged ions.[16]
77
4.2 Results and Discussion
4.2.1 Synthesis and NMR Spectroscopic Characterization
1-(BOhfip2)Fc 2 and 1,1’-(BOhfip2)2Fc 3: The preparation of the
bromoborylferrocene starting materials was realized by electrophilic substitution
of ferrocene with borontribromide according to Siebert et al.[9] First attempts to
synthesize 3 from bromoborylferrocene and HO-hfip alcohol were not promising:
however, we could characterize the desired compound as a byproduct. The main
product of the reaction was ferrocene. We suggest that compound 3 was formed,
but that the B-C bond was cleaved by the HBr released giving ferrocene and the
volatile alkoxybromoborane Br-B(Ohfip)2. To prevent the formation of HBr, the
alcohol was replaced by the alkoxide Li[Ohfip].[17] The consecutively generated
LiBr could be removed and the clean product 2 and 3 were obtained in good yield
(71 % 2; 81 % 3) after recrystallization from hexane or CH2Cl2, respectively
(Scheme 4.1).[18]
Li[Ohfip]; hexane, RT, 2 h- LiBr
B(Ohfip)2
B(Ohfip)2
Fe
0,1
BBr2
BBr2)
Fe
0,1(2, 3)
Scheme 4.1: Synthesis of 2 and 3 from the respective bromoborylferrocene and Li[Ohfip].
The 1H NMR resonances of 2 appear at 5.26 ppm (septet, Ohfip-substituent),
4.64 and 4.46 ppm (two signals of the AA’BB’ spin system), and at 4.23 ppm
(singlet, unsubstituted Cp). The corresponding integral ratios are consistent with
the assigned molecular structure. The 19F NMR spectrum shows a doublet of the
fluorinated alkoxy ligand at −75.4 ppm. The signals of the Ohfip substituents are
shifted (1H: by 0.89 and 0.85 ppm, 19F: by 1.0 and 2.6 ppm) to higher field in both
spectra compared to the alcohol and alkoxide, respectively. The 11B NMR signal
of compound 2 occurs at 31.8 ppm (Δν1/2 = 208 Hz).
Compound 3 exhibits two signal groups, a septet at 5.16 ppm and the multiplets
of the substituted Cp rings at 4.45 and 4.60 ppm. To prove the molecular
structure we consulted selected 2D NMR experiments. The nuclear Overhauser
enhancement (NOE) in the 1H,1H NOESY NMR spectrum (Figure 4.2) between
78
the septet and the signal of the Cp ring at 4.45 proved the small spatial
separation and that both signals belong to one molecule. Furthermore, it allowed
the unambiguous assignment of the signal at 4.45 ppm to the protons of the Cp
ring adjacent to the B substituent. Accordingly, the remaining signal is assigned
to the other protons of the Cp ring.
ppm
4.24.44.64.85.05.25.45.6 ppm
4.2
4.4
4.6
4.8
5.0
5.2
5.4
5.6
Figure 4.2: Section of the 1H,1H–NOESY NMR spectrum (400.17 MHz; mixing time: 0.8 s) of compound 3 in CD2Cl2 at 298K.
Via the cross peaks of the 1H,13C HSQC NMR experiment (appendix, Figure
4.11), it was possible to relate the proton signals to carbon resonances. The
proton of the Ohfip ligand is directly bound to the C atom with the signal at 70.5
ppm. The signals of the ring protons at 4.45 and 4.60 ppm correlate to the C
resonances at 74.7 and 75.0 ppm The signals of the quaternary carbon atom are
identified at shift to higher field, the 2J(1H,13C) cross signals to the ring protons
are observable. In the 19F NMR spectrum, a doublet at −75.5 ppm and in the 11B
NMR spectrum a singlet at 31.2 ppm were observed (Δν1/2 = 365 Hz).
Also, vibrational spectra of 2 and 3 were recorded for completion and show the
expected bands.
79
Oxidation to [1-(BOhfip2)Fc]+[Al(ORF)4]− 4 and [1,1’-(BOhfip2)2)Fc]+[Al(ORF)4]− 5: Ag+[Al(ORF]4]− was applied as the simple to use oxidizing agent of choice
(Scheme 4.2), which was synthesized according to the procedure established in
our group.[14]
B(Ohfip)2
B(Ohfip)2
FeAg+[Al(ORF)4]-; CH2Cl2, RT, 15 min
- Ag
B(Ohfip)2
B(Ohfip)2
Fe [Al(ORF)4]-
0,1 0,1
(4, 5)
Scheme 4.2: Synthesis of 4 and 5 from 2 and 3 with Ag[Al(ORF)4].
The claret-red ferrocenylboranes 2 and 3 were oxidized with Ag+[Al(ORF]4]− to
deep blue ferrocinium derivates. The concomitantly formed elemental Ag was
filtered off, the crude product was washed and recrystallized from CH2Cl2 at 2°C
(89 % 4; 91 % 5). Owing to the paramagnetism of the oxidized compounds, the
NMR spectra were not meaningful; however, vibrational spectra were recorded
for completion and show the expected bands in agreement with the assigned
molecular structures.
Additionally UV/VIS spectra of all compounds were recorded showing the shift of
the absorption maxima from blue (2: 448 nm; 3: 450 nm) to red (4: 624 nm;
5: 630 nm) by oxidation.
4.2.2 X-Ray Crystal Structures
A crystallographic study of single crystals confirmed the proposed connectivity of
the compounds 2-5. The structure of 2 (orthorhombic, Pbca) is shown in Figure
4.3, along with selected metric data.
The dip angle α* of 7.8° is found in the expected range (FcB(Me)2 = 13.0°,
FcB(OH)(Me) = 10.8°, and FcB(OMe)2 = 5.6°).[11] The electron-withdrawing effect
of the polyfluorinated groups on the B-O π-bonding is probably responsible for
the increase of the dip angle α* of 2.2° if compared to FcB(OMe)2. The same
trend is observed when studying the Fe-B distance of 312.6 pm
80
(FcB(Me)2 = 300.8 pm, FcB(OH)(Me) = 305.6 pm, and FcB(OMe)2 = 317.9
pm).[11] The B atom is coordinated trigonal planar, the sum of the angles about B1
is 360°, but both angles around the Cipso-C are significantly different: the C1-B1-
O1 angle is compressed with 115.2° compared to the wide C1-B1-O2 angle of
129.3°. This effect occurs also in FcB(OH)2 (118.2° and 123.9°), but is more
pronounced In ferrocenyl alkoxides like FcB(OMe)2 (e.g., 114.9° and 127.7°).[11,
19] The Ohfip moieties are twisted out of the plane defined by the five Cp carbons
by 15.8°, respectively, as previously observed for ferrocenylboranes.[20] This
pattern can be considered as a sign for a relatively weak Fe-B interaction.
C13
C1
C2C3
C4
B1
C6H1
C8
C7
O1
F1F2
F4
F3
F6
F5 C9
C16
C5
H2Fe1
C11
C10
C15
O2
C14
C12
F7
F12 F11
F10
F9
F8
(2)
Figure 4.3: Molecular structure of 2; thermal ellipsoids are shown at the 50 % probability level. H atoms of the ferrocenyl moiety omitted for clarity. Selected atom distance [pm], angles [°] and interplanar angels [°]: B1-C1 = 153.7(2), B1-O1 = 133.7(2), B1-O2 = 138.1(2), O1-C6 = 140.5(2), O2-C9 = 140.0(1), Fe1-B1 = 312.6(1), Fe1-O1 = 394.6, Fe1-O2 = 393.6; C1-B1-O1 = 115.2(1) , C1-B1-O2 = 129.3(1), O1-B1-O1 = 115.5(1); BO2-Cp: 15.8; COG-Fe1-COG = 176.3 (COG : geometrical centre of the plane defined by the five carbon atoms); α* = 7.8.
Figure 4.4 shows the molecular structure of compound 3 (monoclinic, C2/c)
together with selected structural data. The dip angels of 1,1’-ferrocenylboranes
are usually smaller than the dip angles of the corresponding mono-borylated
ferrocenyles (e.g. FcBBr2/1,1’Fc(BBr2)2 α* = 18.3/9.1° or
81
FcB(OMe)2/1,1’Fc(B(OMe)2)2 α* = 6.1/5.2°).[10, 11, 21] In our case, the dip angle α*
is 8.0° and therefore slightly larger than in compound 2, which could result from
the steric demand of the perfluorinated isopropoxy ligands or packing effects.
Concomitantly, the Fe1-B1 distance is a bit shorter and the boron atoms are
trigonal planar coordinated, with similar bond angels as compound 2. Compared
to the monoborylated compound 2, the Ohfip moieties are even more twisted
(17.2°).
C4
C2
C6
C3
C5 C7
C8
Fe1
C9
C10
C11
C1
C22
C21
C16C15
C14
C13C12
C20
C19
C18
C17
B2
B1
H4
H3
O1
H2
H1
O3
O2O4
(3)
Figure 4.4: Molecular structure of 3; thermal ellipsoids are shown at the 50 % probability level. H atoms of the ferrocenyl moiety omitted for clarity. Selected atom distance [pm], angles [°] and interplanar angels [°]: B1-C1 = 153.1(3), =, B1-O1 = 137.7(3), B1-O2 = 137.4(3), O1-C6 = 140.6(3), O2-C9 = 139.9(3), Fe1-B1 = 309.1; Fe1-O1 = 363.1, Fe1-O2 = 412.0; C1-B1-O1 = 117.0(2), C1-B1-O2 = 128.5(2); BO2-Cp: 17.2; COG-Fe1-COG = 179.9; α* = 8.0.
A section of the molecular structure of compound 4 is shown in Figure 4.5;
relevant metric parameters are included. The non-interacting counterion
[Al(ORF)4]− was omitted for clarity. There are two different motives of 4 in the
asymmetric unit. The metric parameters (the BO2-/Cp-plane angles excluded) do
not significantly distinguish and in the following the average values are
discussed. The dip angle α* amounts 1.6° and is compared to compound 2 (7.8°)
highly compressed. As expected the dip angle decrease and along with them the
Fe-B interaction, since the trivalent cationic Fe atom is not able to contribute as
much electron density to the Lewis acid B centre as the neutral divalent atom in 2
and 3. The boron atom is still trigonal planar coordinated, but the Ohfip moieties
82
are twisted out of the Cp ring plane by 67.8° (molecule A of two in the
asymmetric unit) and by 80.1° (molecule B), which means that the pz-orbital of
the B atom is not able to interact with the Fe-based orbitals or p orbitals of the
Cipso atom. Furthermore, the difference in the values for the twisting angle shows
that the C1-B1 rotation barrier is quite low and that packaging effects take
stronger influence.
C13
C1C2C3
C4B1
C6
H1
C8
C7
O1
F1
F2
F4
F3
F6
F5
C9
C16
C5
H2
Fe1
C11
C10
C15
O2
C14C12
F7
F12
F11
F10
F9
F8
(4)
Figure 4.5: Section of motive A of 4; thermal ellipsoids are shown at the 50 % probability level. H atoms of the ferrocenyl moiety and the [Al(ORF)4]− anion omitted for clarity. Selected atom distance [pm], angles [°] and interplanar angels [°] of motive A: B1-C1 = 155.5(6), B1-O1 = 136.9(5), B1-O2 = 135.7(5), O1-C6 = 141.9(4), O2-C9 = 141.9(4), Fe1-B1 = 321.4, Fe1-O1 = 342.2, Fe1-O2 = 447.6; C1-B1-O1 = 117.5(4) , C1-B1-O2 = 125.9(4) , O1-B1-O1 = 116.6(4); BO2-Cp: 67.8; COG-Fe1-COG = 178.9; α* = 1.5. Selected atom distance [pm], angles [°] and interplanar angels [°] of motive B: B1-C1 = 157.2(6), B1-O1 = 136.7(5), B1-O2 = 134.8(5), O1-C6 = 141.9(4), O2-C9 = 141.6(4), Fe1-B1 = 322.3, Fe1-O1 = 339.0, Fe1-O2 = 450.0; C1-B1-O1 = 117.1(4), C1-B1-O2 = 126.1(4) , O1-B1-O1 = 116.8(4); BO2-Cp: 80.1; COG-Fe1-COG = 179.4; α* = 1.5.
Figure 4.6 shows a section of the molecular structure of compound 5 with
selected metric data. The non-interacting counterion [Al(ORF)4]− was omitted for
clarity. There are two different dip angles α* in the oxidized species of compound
3, in fact 5.4° (at B2) and 0.3° (at B1). Especially, the small dip angle of 0.3° at
B1 indicates that the interaction is almost turned off by oxidation. Furthermore,
this angle is one of the smallest dip angles ever measured in ferrocenylboranes.
Analogous to compound 4 the Ohfip moieties are twisted out of the Cp ring plane
83
by 36.9° at B2 and actually 66.1° at B1. These results, combined with those of
compound 4, show that the Fe-B stabilization is overcompensated by the steric
aspect of the bulky perfluorinated ligands, which prefer to avoid the Cp ring
protons or the second B(Ohfip)2 residue.
C4
C2C6
C3
C5
C7
C8Fe1
C9
C10
C11
C1
C22
C21
C16C15
C14C13
C12
C20
C19
C18C17
B2
B1
H4
H3
O1
H2
H1
O3
O2
O4
(5)
Figure 4.6: Molecular structure of 5; thermal ellipsoids are shown at the 50 % probability level. H atoms of the ferrocenyl moiety and the [Al(ORF)4]− anion omitted for clarity. Selected atom distance [pm], angles [°] and interplanar angels [°]: B1-C1 = 1.575(9), B2-C12 = 1.571(9), B1-O1 = 1.365(8), B1-O2 = 1.354(8), B2-O3 = 1.352(8), B2-O4 = 1.364(8), O1-C6 = 1.412(7), O2-C9 = 1.412(7) , O3-C17 = 1.404(7), O4-C20 = 1.414(7), Fe1-B1 = 3.238; Fe1-B2 = 3.194, Fe1-O1 = 3.454, Fe1-O2 = 4.476, Fe1-O3 = 4.322, Fe1-O4 = 3.518; C1-B1-O1 =117.0(5), C1-B1-O2 = 123.9(6), O1-B1-O2 = 119.0(5), C12-B2-O3 = 126.4(6), C12-B2-O4 = 115.1(6), O3-B2-O4 = 118.5(6); BO2-Cp: 66.1 (at B1), 36.9 (at B2); COG-Fe1-COG = 179.8; α* = 0.3 (at B1), 5.4 (at B2).
By comparing the neutral species 2 and 3 with the oxidized species 4 and 5
(Table 4.1), it follows that in 4 and 5 the C-B distances increase while the B-O
distances decrease, caused by the changed electronic environment. To relieve
the electron deficiency of the B atom in 4 and 5, the B-O π-interaction becomes
more important. The huge difference of the interplanar angles upon oxidation
84
shows that the influence of the sterically demanding Ohfip moieties increases and
therewith overcompensates the Cipso-B π-interaction. The resulting Fe-B
distances and α* angles indicate only a residual and weak Fe-B interaction in 4
and 5. It would be possible that the origin of the twisting results from a Coulomb
interaction between Fe and O (e.g. in 5; dFe1-O1 = 345.4 pm), the affected B-O
bond is not significantly longer, but in contrast the B-O bond, which is facing
away from Fe1, is longer than its counterpart.
Table 4.1: Bond distances d [pm], interplanar angles [°] and dip angles [°] of 2, 3, 4 and 5 (average values in parentheses).
dCB dBO dFeB BO2-/Cp-plane α*
2 153.7 (153.7) 137.7-138.1 (137.9) 312.6 15.8 7.8
3 153.1 (153.)1 137.4-137.6 (137.5) 309.1 17.2 8.0
4 155.5-157.2 (156.4) 134.8-136.9 (136.0) 342.2 (A),
339.0 (B)
67.8 (A) , 80.1 (B) 1.5 (A), 1.6 (B)
5 157.1-157.4 (157.3) 135.3-136.5 (135.9) 323.8 (B1),
319.4 (B2)
66.1 (B1), 36.9 (B2) 0.3 (B1), 5.4 (B2)
4.2.3 Cyclo Voltammetric Characterization
To further investigate the electron-withdrawing potential of the B(Ohfip)2 group,
cyclic voltammetric experiments were run on the compounds 2 and 3. The cyclic
voltametric measurements were performed with a platinum ultra micro electrode
(UME; r = 5 μm). Before measurement the UME was polished with different
diamond pastes with smallest particle size of 0.25 μm. The effective radius (reff)
was determined by measuring the steady state current (Iss) in an electrolyte
solution of 1 mM ferrocene, 100 mM [nBu4N]+[ClO4]− in acetonitrile to reff 7.2 μm.
Platinum wires were used as counter electrode and quasireference electrode,
respectively, and the corresponding redox potentials are referenced against the
Fc/Fc+ couple. Measurements were performed with 100 mM solutions of
[nBu4N]+[Al(OC(CF3)3)4]− as supporting electrolyte and with 1 mM solutions of
compound 2 and 3 in CH2Cl2. The scan rate was 10 mV s−1. After the
measurements of the samples Fc was added to the electrolyte solution and for
referencing the cyclic voltammetric measurements were repeated.
85
Figure 4.7 and Figure 4.8 show the cyclic voltammograms of 2 and 3. The curve
progression of compound 2 (Figure 4.7) fits well with the expected progression
apart from a slightly too steep rise between 0.2 and 0.4 V, which can be
explained by small impurities.
-1.0 -0.5 0.0 0.5 1.0 1.5-1.0x10-9
0.0
1.0x10-9
2.0x10-9
3.0x10-9
4.0x10-9
5.0x10-9
I [A]
E [V] vs. q-Pt
Compound 2
Figure 4.7: Cyclic voltammograms of 2 (blue). Measurements were carried out in CH2Cl2 solutions with [nBu4N]+[Al(OC(CF3)3)4]− as supporting electrolyte at a scan rate of 10 mV s−1.
The curve progression of compound 3 (Figure 4.8) is surprisingly flat. This also
may arise from impurities of the sample of compound 3, which can lead to
electrode fowling. However, if it is assumed that the sample is free of detectable
impurities the diffusion coefficient D and the rate constant of the heterogeneous
electron transfer k0 of 3 are significantly smaller than expected. Therefore, the
results shown in Table 4.2 should be interpreted with care.
86
-1.0 -0.5 0.0 0.5 1.0 1.5-1.0x10-10
0.0
1.0x10-10
2.0x10-10
3.0x10-10I [
A]
E [V] vs q-Pt
Compound 3
Figure 4.8: Cyclic voltammograms of 3 (red). Measurements were carried out in CH2Cl2 solutions with [nBu4N]+[Al(OC(CF3)3)4]− as supporting electrolyte at a scan rate of 10 mV s−1.
By simulations with the program DigiElch® (Figure 4.9 and Figure 4.10) the formal
potentials E0’, the rate constant of the heterogeneous electron transfer k0, the
transfer coefficient α and the diffusion coefficient D were calculated (Table 4.2).
The formal potential E0’ is defined by (E0: redox potential; z: number of electrons;
γ: activity coefficients oxidized and reduced species, respectively):
It can be postulated that γox ≈ γred, which in turn leads to the assumption that
E0’ ≈ E0. Additionally it can be assumed that Dox = Dred (as done for the
simulation), and thereby E0’ = E1/2.
87
-1.0 -0.5 0.0 0.5 1.0 1.5-2.0x10-9
0.0
2.0x10-9
4.0x10-9
6.0x10-9
8.0x10-9
1.0x10-8
1.2x10-8
1.4x10-8I [
A]
E [V] vs q-Pt
Compound 2 and Fc Simulated
Figure 4.9: Cyclic voltammograms of 2 (blue) after adding Fc and simulated curves (black).
-1.0 -0.5 0.0 0.5 1.0 1.5
0.0
2.0x10-9
4.0x10-9
6.0x10-9
8.0x10-9
1.0x10-8
1.2x10-8
I [A]
E [V] vs. q-Pt
Compound 3 and Fc Simulated
Figure 4.10: Cyclic voltammograms of 3 (red) after adding Fc and simulated curves (black).
88
Table 4.2: Thermodynamic and kinetic data of the electrode reaction of compound 2 and 3 in CH2Cl2 obtained from the simulation curves shown in Figure 4.9 and Figure 4.10.
Fc Compound 2 Compound 3
E0’ [V] 0 0.320 0.747
k0 [cm s−1] 1e−1 2e−2 2e−4
α 0.58 0.58 0.58
D [cm2 s−1] [a]1.67e−5 1.23e−5 1.07e−6
[a] taken from literature.[22]
The formal potentials E0’ of compound 2 and 3 in CH2Cl2 amount to +0.32 and
+0.75 V versus Fc/Fc+, respectively. The anodic shift in 2 can be contributed to
the π-electron-withdrawing effect of the B(Ohfip)2 group and is nearly more than
doubled in 3, due to the introduction of a second B(Ohfip)2 group at a formerly
unsubstituted Cp ring. Furthermore, in the same way the oxidation potential
increases, since the Fe atom is more difficult to oxidize due to electron-
withdrawing effect. This is in congruence to the values for FcB(C6F5)2 (E1/2 =
0.450 V versus Fc/Fc+), FcB(OMe)2 (E0’ = 0.09 V versus Fc/Fc+) and
1,1’Fc(B(OMe)2)2 (E0’ = 0.26 V versus Fc/Fc+).[11, 15] Since the B(C5F5)2 residue is
a stronger electron-withdrawing residue than the B(Ohfip)2 group, the anodic shift
is more pronounced in FcB(C6F5)2 than in 2. In contrast, the B(OMe)2 ligand does
not contain any electro negative fluorine atoms and thus the E0’ value lies closer
to Fc.
Concerning the rate constant k0 of Fc the relatively low value compared to data
from literature (k0 > 6)[23] may arise from impurities as mentioned above. Possibly
the k0 values of 2 and 3 are affected by the same uncertainty as well.
Nevertheless, it can be seen that the electron transfer slows down by increasing
the number of substituents and reflects their steric influence. The transfer
coefficient specifies the location of the activated complex within the
electrochemical double layer and can take values between 0 and 1. Usually
values between 0.3 and 0.7 are observed, therefore the obtained values are
reasonable. From the Iss value the diffusion coefficient of the compounds were
determined to D(2) = 1.23e−5 and D(3) = 1.07e−6 cm2 s−1. The D(3) value would
be expected to be only slightly lower than the D(2) value, however the
discrepancy of this magnitude indicates that it is likely that electron fowling
occurred.
89
4.2.4 Quantum Chemical Investigations
The CIA, HIA, FIA and MIA values of the neutral compounds 2 and 3 have been
determined accordingly to the method of Chapter 3. Since there is a rotation
barrier around the Cipso-B bond, different stable isomers are possible. Hence,
herein we started our geometry optimization with three different start geometries
(Scheme 4.3) and discussed the converged structures with the lowest SCF
energies. Additionally, we used bridged start geometries for the CIA and FIA of
the bifunctional Lewis acid 3 to control if a bridged structure occurs in those
anions. For calculating the Lewis acids the crystal structure was used as starting
geometry. The complete energetic data can be found in the appendix of the
chapter.
B
Fe
0,1(hfipO)2B
Ohfip
OhfipX
B
Fe
0,1(hfipO)2B
X
Ohfip
hfipO
hfipO
B
Fe
0,1(hfipO)2B
Ohfip
XhfipO
B
Fe
B
X'
Ohfip
Ohfip
hfipO
"in" "out"
"side-on" "bridged"
Scheme 4.3: Starting geometries to calculate the CIA, HIA, FIA and MIA values of compound 2 and 3. (X = Cl, H, F, Me; X’ = Cl, F).
However, no bridged motive was obtained in the optimized structures. Since the
((hfipO)2B)Cp-Fe-Cp(B(Ohfip)2) axis is rotatable, the bridged start geometries
converged to structures without staggered B(Ohfip)2 moieties.
Table 4.3 shows the Lewis affinity values of compound 2 and 3 of their most
stable calculated isomers.
90
Table 4.3: CIA, HIA, FIA and MIA values of 2 and 3. The conformation of the starting geometry is given in parenthesis.
CIA
[kJ mol−1]
HIA
[kJ mol−1]
FIA
[kJ mol−1]
MIA
[kJ mol−1]
2 168 (a) 352 (b) 377 (b) 380 (a)
3 212 (d) 383 (b) 408 (b) 403 (a)
As excepted the Lewis affinity values raise upon adding a second B(Ohfip)2
moiety, since the introduction of a further electron poor boron atom increases the
electron deficiency of the molecule. Furthermore, it is a sign that the two boron
atoms in 3 communicate electronically via the Fe atom. In compound 2 only one
boron atom withdraws electron density from the Fe atom. In contrast, in
compound 3 two boron atoms compete for the same amount of electron density,
which results in two more electron deficient boron atoms and a higher Lewis
acidity. Consequently, both boron atoms recognize the presence and type of
electronic structure of each other.
To estimate the Lewis acidity of the cationic Lewis acids 4 and 5 we tried to
calculate the gas phase values using the same procedure as above. To reduce
the cation/anion interaction and to gain comparable values, we wanted to transfer
the Lewis acids 2, 3, 4 and 5 and their respective most stable Lewis acid/base
complexes to solution (using the COSMO solvation model). As imaginary solvent
CH2Cl2 seemed to be a good choice since it was used for FIA calculations of
phosphenium ions.[24]
By using this method, preliminary calculations of the FIA of 5 in CH2Cl2, figure out
that oxidation increases the FIA for about 100 kJ mol−1. However, to charge
separation compounds it is unlikely that the results are reliable, as DFT methods
tend to smear charges and the orbitals might mix unphysical.
91
4.3 Conclusion
Herein the straightforward synthesis and characterization of [1-(BOhfip2)Fc] 2 and
[1,1’-(BOhfip)2)2Fc] 3 and the subsequent oxidation of 2 and 3 with Ag+[Al(ORF)4]−
to the ferrocinium derivates [1-(BOhfip2)Fc]+[Al(ORF)4]− 4 and
[1,1’-(BOhfip2)2)Fc]+[Al(ORF)4]− 5 was reported. Furthermore, the crystal
structures of 2 and 3 and even the rare crystal structures of the oxidized
ferrocenylboranes 4 and 5 were obtained. It could be shown that the dip angle α*
strongly decrease upon oxidation and the Fe-B interaction is enfeebled. In
addition, the Ohfip moieties twist out of the Cp plane to minimize steric
interactions with the Cp ring protons or the second B(ORF)2 residue. Cyclic
voltametric measurements of 2 and 3 were performed and the thermodynamic
and kinetic data obtained. The anodic shift of 2 (E0’ = +0.32 V versus Fc+/Fc) and
3 (E0’ = +0.75 V versus Fc+/Fc) reflect very well the introduction of electron-
withdrawing B(Ohfip)2 residues. Furthermore, the CIA, HIA, FIA and MIA values
of the neutral compounds 2 and 3 have been determined in consideration of
different rotamers. Also here the introduction of B(Ohfip)2 groups can be
observed in the higher affinity values of compound 3 compared to 2.
92
4.4 Experimental Section
4.4.1 Theoretical Methods
All geometries were optimized at the (RI-)BP86/SV(P) level[25] using
TURBOMOLE 6.4.[26] Vibrational frequencies were calculated with the AOFORCE
module[27] and all structures represented true minima without imaginary
frequencies. The reference values for the Lewis acidity scale were taken from
Chapter 3.
4.4.2 Synthesis and Characterization
Techniques and Instruments: Due to the air and moisture sensitivity of most
materials, all manipulations were undertaken with vacuum and Schlenk
techniques as well as in a glove box with an argon atmosphere (H2O and O2 < 1
ppm). The solvents were dried by using conventional drying agents and directly
distilled. NMR spectra were obtained in d2-methylene chloride at room
temperature on a BRUKER AVANCE II+ 400 spectrometer. 1H and 13C chemical
shifts are given with respect to TMS, 19F NMR spectra to fluorotrichloromethane
and 11B NMR spectra to the boron-trifluoride-diethyl-ether-complex. IR spectra
were recorded on a ZnSe or Diamond ATR unit on a Nicolet Magna IR
Spectrometer with 1 cm–1 resolution or on a Diamond ATR unit (200 - 30000
cm–1) on a BRUKER alpha FT-IR Spectrometer with a resolution of 4 cm–1.
Raman spectra were obtained on a BRUKER VERTEX 70 spectrometer with a
BRUKER RAM II Raman unit in sealed melting point capillaries. Data collections
for X-ray structure determinations were performed on a Rigaku Spider image
plate or a BRUKER APEX II Quazar CCD diffractometer at 100K and 110K,
respectively, with Mo radiation. The single crystals were mounted in
perfluoroether oil on a MiTeGen MicromountTM. UV/VIS spectra were recorded on
an Evolution 600 spectrometer with a DRA-EV-600 integration sphere (75 mm
diameter). The samples were measured in a specimen holder with teflon ground
and a quartz window in a range of 300 – 750 nm. Cyclic voltametric
measurements were performed with a platinum ultra micro electrode (UME; r = 5
μm). Before measurement the UME was polished with different diamond pastes
93
with smallest particle size of 0.25 μm. The effective radius (reff) was determined
by measuring the steady state current (Iss) in an electrolyte solution of 1 mM
ferrocene, 100 mM [nBu4N]+[ClO4]− in acetonitrile to reff 7.2 μm. Platinum wire
was used as counter electrode and quasireference electrode, respectively, and
the potential values herein are given relative to the Fc/Fc+ couple. Measurement
were performed with 100 mM solutions of [nBu4N]+[Al(OC(CF3)3)4]− as the
supporting electrolyte and with 1 mM solutions of compound 2 and 3 in CH2Cl2.
The scan rate was 10 mV s−1. After the measurements of the samples Fc was
added to the electrolyte solution and for referencing the cyclic voltametric
measurements were repeated. Thermodynamic and kinetic data were obtained
by simulations with the program DigiElch®. Dibromoborylferrocene,
1,1’-bis(dibromoboryl)ferrocene, Ag+[Al(ORF)4]– and LiOhfip were synthesized
according to the established procedures[9, 14, 17]
Preparation of compound 2: Dibromoborylferrocene (1.342 g, 3.78 mmol) and
LiOhfip (1.314 g, 7.55 mmol) were placed on one side of a special two-flask
fritplate vessel with J. Young valves and dissolved in hexane (70 ml). An orange
solution over colorless precipitate formed and was left stirring for two hours. The
mixture was filtered to the other side of the vessel, the solvent was removed in
vacuo and the solid was placed in clean two-flask fritplate vessel with J. Young
valves and dissolved in hexane (50 ml). After recrystallization of the product at –
40°C, the excess filtrate was poured to the other side of the vessel and the
residual solvent was removed in vacuo. Compound 1 was obtained on the
starting side of the vessel (1.419 g, 2.68 mmol, 71%). 1H NMR (400.17 MHz):
δ = 4.21 (s, 5H, Cp), 4.44 (m, 4H BCCH), 4.62 (m, 4H, BCCHCH), 5.23 (sept.,
4H, CB(Ohfip)2), 5.32 (t, 1H, CHDCl2). 19F (376.54 MHz): δ = –75.4 (d, 12F CF3).
11B (128.39 MHz): δ = 31.5 (br). IR (Diamond ATR, corrected): ν~ = 483 (s), 503
(w), 522 (vw), 627 (vw), 674 (w), 687 (s), 697 (w), 742 (w), 754 (vw), 818 (w), 834
(m), 865 (m), 872 (m), 891 (w), 903 (m), 1001 (w), 1024 (w), 1038 (w), 1066 (w),
1103 (vs), 1192 (vs), 1218 (s), 1258 (s), 1281 (m), 1305 (m), 1337 (w), 1365 (w),
1381 (w), 1461 (w), 2958 (vw), 3108 (vw), 3203 (vw) cm–1. Raman: ν~ = 75, 131,
223, 256, 277, 316, 389, 433, 485, 597, 628, 676, 688, 700, 754, 836, 865, 892,
1002, 1024, 1041, 1054, 1068, 1081, 1106, 1180, 1197, 1211, 1290, 1322, 1353,
1380, 1398, 1411, 1461, 2960, 2981, 3108, 3118 cm–1.
94
Preparation of compound 3: 1,1’-Bis(dibromoboryl)ferrocene (1.440 g, 2.74
mmol) and LiOhfip (2.011 g, 11.56 mmol) were mixed in the glovebox, weighed
into one side of a special two-flask fritplate vessel with J. Young valves and
dissolved in CS2 (80 ml). The mixture was stirred at room temperature for two
hours and then washed two times with CS2. The residual solvent was removed
and the product extracted with CH2Cl2. Upon removal of the solvent from this
extracted filtrate, microcrystalline compound 2 was obtained in good yield (1.962
g, 2.24 mmol, 81%). Crystals suitable for XRD were obtained by recrystallization
from CH2Cl2 at 2°C. 1H NMR (400.17 MHz): δ = 4.47 (m, 4H, BCCH), 4.62 (m,
4H, BCCHCH), 5.17 (sept., 4H, CB(Ohfip)2), 5.32 (t, CHDCl2). 19F (376.54 MHz):
δ = –75.1 (d, 24F, CF3). 11B (128.39 MHz): δ = 31.4 (br). 13C (100.62 MHz):
δ = 53.5 (m, CD2Cl2), 57.4 (m, BCCH2CH2), 70.5 (m, OCH(CF3)2), 74.7 (m,
BCCH2CH2), 75.0 (m, BCCH2CH2), 120.8 (m, OCH(CF3)2). IR (ZnSe ATR,
corrected): ν~ = 673 (m), 690 (vs), 741 (m), 810 (w), 820 (w), 844 (w), 864 (w),
874 (m), 904 (w), 1024 (w), 1038 (w), 1103 (vs), 1196 (vs), 1226 (s), 1263 (vs),
1284 (s), 1325 (w), 1336 (m), 1363 (w), 1380 (m), 1462 (w), 2852 (vw), 2920
(vw), 2964 (vw) cm–1. Raman: ν~ = 131, 206, 298, 325, 333, 345, 401, 451, 474,
626, 673, 700, 755, 838, 865, 875 892, 1024, 1039, 1064, 1081, 1108, 1182,
1197, 1276, 1290, 1311, 1326, 1340, 1382, 1398, 1459, 1471, 1496, 2971, 3122
cm–1.
Preparation of compound 4: Compound 2 (0.106 g, 0.200 mmol) and
Ag+[Al(ORF)4]- (0.213 g, 0.200 mol)) were mixed in the glovebox and weighed into
one side of a special two-flask fritplate vessel with J. Young valves and dissolved
in CH2Cl2 (6 ml). The blue solution was stored at −20°C for several months until
suitable crystals for XRD were obtained. The solvent was removed and 5 (0.275
g, 0.178 mmol, 89%) was characterized. IR (Diamond ATR, corrected): ν~ = 443
(vw), 537 (vw), 561 (vw), 693 (w), 725(m), 742 (vw), 755 (vw), 833 (vw), 857
(vw), 874 (vw), 907 (vw), 968 (s), 1101 (w), 1111 (m), 1170 (m), 1209 (vs), 1238
(s), 1267 (s), 1292 (m), 1353 (vw), 1379 (w), 1393 (vw), 1424 (vw), 1468 (vw),
3126 (vw), 3131 (vw) cm–1. Raman: ν~ = 167, 268, 308, 368, 538, 555, 746, 800,
1002, 1072, 1116, 1185, 1266, 1488, 1496, 1990, 3128 cm–1.
95
Preparation of compound 5: Compound 3 (0.400 g, 0.458 mmol) and
Ag+[Al(ORF)4]- (0.488 g, 0.458 mol)) were mixed in the glovebox and weighed into
one side of a special two-flask fritplate vessel with J. Young valves and dissolved
in CH2Cl2 (15 ml). The blue solution was stirred at room temperature for 15
minutes filtered and washed once with CH2Cl2 (2 ml). The solvent was removed
and compound 3 was obtained in good yield (0.771 g, 0.421 mmol, 91%).
Crystals were obtained by recrystallization from CH2Cl2 at 2°C. IR (Diamond
ATR, corrected): ν~ = 445 (m), 481 (vw), 536 (m), 560 (w), 692 (s), 726 (vs), 744
(w), 755 (w), 828 (w), 867 (s), 906 (m), 968 (vs), 1040 (vw), 1094 (s), 1111 (vs),
1197 (vs), 1240 (s), 1265 (s), 1352 (w), 1373 (m), 1421 (vs), 1468 (vw), 2239
(vw), 2361 (vw), 2857 (vw), 2920 (vw) cm–1. Raman: ν~ = 79, 167, 235, 293, 322,
500, 560, 746, 797, 856, 885, 1063, 1104, 1179, 1278, 1386, 1459, 1488, 1504,
1990, 2063, 2094, 2329, 2688, 2758, 2846, 2938, 3125 cm–1.
96
4.5 References
[1] C. Bresner, S. Aldridge, I. A. Fallis, C. Jones, L. L. Ooi, Angewandte Chemie-International Edition 2005, 44, 3606; C. Dusemund, K. Sandanayake, S. Shinkai, Journal of the Chemical Society-Chemical Communications 1995, 333; S. Aldridge, C. Bresner, I. A. Fallis, S. J. Coles, M. B. Hursthouse, Chemical Communications 2002, 740; C. Bresner, J. K. Day, N. D. Coombs, I. A. Fallis, S. Aldridge, S. J. Coles, M. B. Hursthouse, Dalton Transactions 2006, 3660.
[2] F. Jäkle, K. Polborn, M. Wagner, Chemische Berichte 1996, 129, 603; F. F. de Biani, F. Jäkle, M. Spiegler, M. Wagner, P. Zanello, Inorganic Chemistry 1997, 36, 2103; E. Herdtweck, F. Peters, W. Scherer, M. Wagner, Polyhedron 1998, 17, 1149; S. L. Guo, F. Peters, F. F. de Biani, J. W. Bats, E. Herdtweck, P. Zanello, M. Wagner, Inorganic Chemistry 2001, 40, 4928.
[3] M. Fontani, F. Peters, W. Scherer, W. Wachter, M. Wagner, P. Zanello, European Journal of Inorganic Chemistry 1998, 1453; M. Grosche, E. Herdtweck, F. Peters, M. Wagner, Organometallics 1999, 18, 4669; R. E. Dinnebier, M. Wagner, F. Peters, K. Shankland, W. I. F. David, Zeitschrift Fur Anorganische Und Allgemeine Chemie 2000, 626, 1400.
[4] M. Herberhold, U. Dorfler, W. Milius, B. Wrackmeyer, Journal of Organometallic Chemistry 1995, 492, 59.
[5] F. F. de Biani, T. Gmeinwieser, E. Herdtweck, F. Jäkle, F. Laschi, M. Wagner, P. Zanello, Organometallics 1997, 16, 4776; L. Ding, K. B. Ma, F. F. de Biani, M. Bolte, P. Zanello, M. Wagner, Organometallics 2001, 20, 1041; K. B. Ma, F. F. de Biani, M. Bolte, P. Zanello, M. Wagner, Organometallics 2002, 21, 3979.
[6] M. Scheibitz, R. F. Winter, M. Bolte, H. W. Lerner, M. Wagner, Angewandte Chemie-International Edition 2003, 42, 924.
[7] W. E. Piers, G. J. Irvine, V. C. Williams, European Journal of Inorganic Chemistry 2000, 2131; B. E. Carpenter, W. E. Piers, R. McDonald, Canadian Journal of Chemistry-Revue Canadienne De Chimie 2001, 79, 291; R. Boshra, A. Sundararaman, L. N. Zakharov, C. D. Incarvito, A. L. Rheingold, F. Jäkle, Chemistry-a European Journal 2005, 11, 2810.
[8] E. Herdtweck, F. Jäkle, G. Opromolla, M. Spiegler, M. Wagner, P. Zanello, Organometallics 1996, 15, 5524; F. Jäkle, T. Priermeier, M. Wagner, Journal of the Chemical Society-Chemical Communications 1995, 1765; F. Jäkle, T. Priermeier, M. Wagner, Organometallics 1996, 15, 2033; F. Jäkle, M. Mattner, T. Priermeier, M. Wagner, Journal of Organometallic Chemistry 1995, 502, 123.
[9] W. Ruf, T. Renk, W. Siebert, Zeitschrift Fur Naturforschung Section B-a Journal of Chemical Sciences 1976, 31, 1028; W. Ruf, M. Fueller, W. Siebert, Journal of Organometallic Chemistry 1974, 64, 45; T. Renk, W. Ruf, W. Siebert, Journal of Organometallic Chemistry 1976, 120, 1.
[10] A. Appel, F. Jäkle, T. Priermeier, R. Schmid, M. Wagner, Organometallics 1996, 15, 1188.
[11] M. Scheibitz, M. Bolte, J. W. Bats, H. W. Lerner, I. Nowik, R. H. Herber, A. Krapp, M. Lein, M. C. Holthausen, M. Wagner, Chemistry-a European Journal 2005, 11, 584.
97
[12] L. Kaufmann, H. Vitze, M. Bolte, H. W. Lerner, M. Wagner, Organometallics 2008, 27, 6215.
[13] A. Meller, Wojnowsk.M, Monatshefte Fur Chemie 1969, 100, 1489; L. O. Müller, D. Himmel, J. Stauffer, G. Steinfeld, J. Slattery, G. Santiso-Quinones, V. Brecht, I. Krossing, Angewandte Chemie-International Edition 2008, 47, 7659; A. Kraft, J. Beck, I. Krossing, Chemistry-a European Journal 2011, 17, 12975; M. Kuprat, R. Kuzora, M. Lehmann, A. Schulz, A. Villinger, R. Wustrack, Journal of Organometallic Chemistry 2010, 695, 1006.
[14] I. Krossing, Chemistry-a European Journal 2001, 7, 490. [15] B. E. Carpenter, W. E. Piers, M. Parvez, G. P. A. Yap, S. J. Rettig,
Canadian Journal of Chemistry-Revue Canadienne De Chimie 2001, 79, 857.
[16] I. Krossing, I. Raabe, Angewandte Chemie-International Edition 2004, 43, 2066.
[17] A. Reisinger, N. Trapp, I. Krossing, Organometallics 2007, 26, 2096. [18] Preliminary investigations were performed and the crystal structure of 3
were obtained during my diploma thesis at the Universität Freiburg 2010. [19] C. Bresner, S. Aldridge, I. A. Fallis, L. L. Ooi, Acta Crystallographica
Section E-Structure Reports Online 2004, 60, M441. [20] H. Braunschweig, F. M. Breitling, K. Kraft, M. Kraft, F. Seeler, S. Stellwag,
K. Radacki, Zeitschrift Für Anorganische Und Allgemeine Chemie 2006, 632, 269.
[21] B. Wrackmeyer, U. Dörfler, W. Milius, M. Herberhold, Polyhedron 1995, 14, 1425.
[22] N. G. Tsierkezos, Journal of Solution Chemistry 2007, 36, 289. [23] A. M. Bond, T. L. E. Henderson, D. R. Mann, T. F. Mann, W. Thormann, C.
G. Zoski, Analytical Chemistry 1988, 60, 1878. [24] J. M. Slattery, S. Hussein, Dalton Transactions 2012, 41, 1808. [25] A. D. Becke, Physical Review A 1988, 38, 3098; J. P. Perdew, Physical
Review B 1986, 33, 8822; J. P. Perdew, Physical Review B 1986, 34, 7406.
[26] A. Schäfer, H. Horn, R. Ahlrichs, The Journal of Chemical Physics 1992, 97, 2571.
[27] P. Deglmann, F. Furche, Journal of Chemical Physics 2002, 117, 9535; P. Deglmann, F. Furche, R. Ahlrichs, Chemical Physics Letters 2002, 362, 511.
98
99
4.6 Appendix
Legend to the following table and calculations herein
BP86 BP86 density functional
h Hartree, 1H = 2625.5 kJ mol−1
Evrt Sum of vibrational, rotational and translational
energy including ZPE (= FreeH energy)
ZPE Zero point vibrational energy
S Entropy
H Enthalpy
G Gibbs Energy
Table 4.4: Summary of the energetic data belonging to compound 2. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound Rotamer (RI-)BP86/SV(P)
energy [h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
CIA
[kJ mol−1]
HIA
[kJ mol−1]
MIA
[kJ mol−1]
FIA
[kJ mol−1]
FcB(Ohfip)2 2 −3252.297269 767.05 0.84170 −8538169.473 −8538560.118
[FcB(Ohfip)2Cl]− a −3712.518644 767.82 0.88506 −9746484.526 −9746742.560 168
[FcB(Ohfip)2H]− a −3252.942778 783.46 0.85906 −8539847.854 −8540103.136 348
[FcB(Ohfip)2Me]− a −3292.229360 858.38 0.87993 −8642920.247 −8642920.247 376
[FcB(Ohfip)2F]− a −3352.193240 770.35 0.86716 −8800444.046 −8800707.926 380
[FcB(Ohfip)2Cl]− b −3712.517126 767.50 0.87604 −9746480.861 −9746743.206 164
[FcB(Ohfip)2H]− b −3252.944237 782.91 0.85696 −8539852.236 −8540108.999 352
[FcB(Ohfip)2Me]− b −3292.229927 858.52 0.86545 −8642921.596 −8642921.596 377
[FcB(Ohfip)2F]− b −3352.191076 770.35 0.85622 −8800438.362 −8800699.554 374
[FcB(Ohfip)2Cl]− c −3712.512801 767.44 0.88400 −9746469.565 −9746855.696 153
[FcB(Ohfip)2H]− c −3252.937092 783.81 0.84389 −8539832.575 −8539832.575 333
[FcB(Ohfip)2Me]− c −3292.225508 857.94 0.87991 −8642910.575 −8643161.528 366
[FcB(Ohfip)2F]− c −3352.187621 770.34 0.86119 −8800429.301 −8800692.865 365
100
101
Table 4.5: Summary of the energetic data belonging to compound 3. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values.
Compound Rotamer (RI-)BP86/SV(P)
energy [h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
CIA
[kJ mol−1]
HIA
[kJ mol−1]
MIA
[kJ mol−1]
FIA
[kJ mol−1]
1.1’Fc(B(Ohfip)2)2 3 −4853.967840 1079.56 1.30994 −12743059.064 −12743154.099
[1.1’Fc(B(Ohfip)2)2Cl]− a −5314.198247 1080.89 1.33009 −13951364.516 −13951455.787 158
[1.1’Fc(B(Ohfip)2)2H]− a −4854.623088 1096.72 1.30333 −12744737.486 −12744737.486 348
[1.1’Fc(B(Ohfip)2)2Me]− a −4893.910482 1171.12 1.34612 −12847809.138 −12848199.697 375
[1.1’Fc(B(Ohfip)2)2F]− a −4953.872760 1083.47 1.31880 −13005333.551 −13005725.266 380
[1.1’Fc(B(Ohfip)2)2Cl]− b −5314.196467 1080.09 1.32468 −13951393.397 −13951788.615 187
[1.1’Fc(B(Ohfip)2)2H]− b −4854.626760 1096.02 1.30581 −12744772.604 −12745160.137 383
[1.1’Fc(B(Ohfip)2)2Me]− b −4893.912281 1171.15 1.32491 −12847842.004 −12847842.004 408
[1.1’Fc(B(Ohfip)2)2F]− b −4953.870844 1082.81 1.31489 −13005352.150 −13005747.103 399
[1.1’Fc(B(Ohfip)2)2Cl]− c −5314.185756 1080.86 1.31382 −13951397.271 −13951792.292 191
[1.1’Fc(B(Ohfip)2)2H]− c −4854.613544 1096.44 1.29666 −12744762.266 −12745154.300 373
[1.1’Fc(B(Ohfip)2)2Me]− c −4893.899820 1171.30 1.32557 −12847837.310 −12847837.310 403
[1.1’Fc(B(Ohfip)2)2F]− c −4953.863821 1082.97 1.29979 −13005356.521 −13005753.087 403
[1.1’Fc(B(Ohfip)2)2Cl]− d −5314.206076 1080.61 1.31023 −13951418.104 −13951819.450 212
[1.1’Fc(B(Ohfip)2)2F]− d −4953.867270 1081.79 1.29509 −13005343.786 −13005736.987 390
ppm
4.34.44.54.64.74.84.95.05.15.25.35.45.55.6 ppm
45
50
55
60
65
70
75
80
Figure 4.11: Section of the 1H,13C-HSQC NMR spectrum (400.17 MHz) of compound 3 in CD2Cl2 at 298 K
ppm
4.04.24.44.64.85.05.25.45.65.8 ppm
50
60
70
80
90
100
110
120
130
140
Figure 4.12: Section of the 1H,13C-HMBC NMR spectrum (400.17 MHz) of compound 3 in CD2Cl2 at 298 K.
102
5 Approaching New Boron-Based Weakly Coordinating Anions (WCAs)
Parts of the syntheses in this chapter were performed by Anne Asmacher as part
of her bachelor thesis under my supervision. The technical data for all executed
calculations is deposited on the file server of the group of Prof. Krossing (path:
public\Ehemalige\Hannes Böhrer 2014\SI_Disseratation\SI_Tech_Chapt5.pdf).
5.1 Introduction
Weakly coordinating anions (WCAs) have become indispensable in chemistry
and found widespread applications in stabilization of reactive cations,
polymerization reactions, ionic liquids and in battery technology.[1, 2] Since the
[BF4]− anion was one of the first anions that was supposed to be weakly
coordinating,[3] huge progress has been made on the field of borate based WCAs.
Hence, the ability of anions to coordinate weakly, benefits from the volume of the
anion and thereby the size of the ligands. Thus, it was just a logical consequence
to replace the fluoride ligands by the sterically more demanding phenyl residues.
However the resulting [BPh4]− anion was still relatively strongly coordinating and
the B-C bond prone to hydrolysis and cleavage under acidic conditions.[4] To
further optimize this type of borate based WCAs the 3 and 5 positions of the
phenyl ring were equipped with CF3 groups or fluorinated phenyl residues were
used, to obtain [B(ArF)4]− (ArF = C6H3-3,5(CF3)2)[5] or [B(C6F5)4]−.[6] The electron-
withdrawing effect of the CF3 groups and fluoride stabilizes the ions against
electrophilic attack and increase their solubility.[7] Tetrakis(perfluorophenyl)borate
is used e.g. as component in supporting electrolyte salts for anodic processes,
since the weakly coordinating abilities provide a different medium in which to
generate positively charged electrolysis products.[8] Those cationic products may
be prone to nucleophilic attack by the conventionally used supporting electrolyte
anions, like [BF4]−.[9] Furthermore [B(C6F5)4]− is utilized to stabilize low-coordinate
tin and lead cations[10] or as anion in a borenium salt, which is able to activate
hydrogen and to hydrogenate imines and enamines.[11] The ArF residues in
[B(ArF)4]− and the C6F5 groups of [B(C6F5)4]− were modified with larger
103
perfluoroalkylgroups. This induced high synthetic effort, but only small
improvements were gained compared to the well-investigated
[B(C6H3-3,5(CF3)2)4]− and [B(C6F5)4]− anions.[2] Based on
tetrakis(perfluorophenyl)borate novel WCAs have been designed, like
[B(CN•B(C6F5)3)4]−, possessing CN linkers between a central boron and the
boron atom of a B(C6F5)3 group.[12] Another method to further reduce the
coordinating ability was to delocalize the negative charge over two boron
atoms,[13] resulting in highly potent WCAs like [(F5C6)3B(μ-C3N2H3)B(C6F5)3]−.[14]
Since the requirements for a WCA depend on the desired application, a large
pool of different WCAs is necessary for modern chemistry to choose the suitable
WCA from. Herein we focus on cost-effective, fast and simple syntheses of
heteroleptic WCAs. Our goal is to investigate the synthesis of WCAs, based on
increasingly bulky fluorinated alcohols to give [(RFO)2B(C6F5)2]− (ORF =
OC(CF3)3, OCH(CF3)2, OCH2CF3; Figure 5.1).
Synthesis Routes: Wagner et al. introduced a one-pot synthesis of the
[H2B(C6F5)2]− anion 1 from C6F5MgBr/BH3•SMe2,[15] which represents our basic
module herein to design the novel WCAs. Furthermore it is possible to generate
HB(C6F5)2 in situ from compound 1 and to trap it with SMe2,[15] which can be a
useful tool to synthesize Lewis acids or WCAs. To get an overview over the
properties of the mentioned WCAs we performed preliminary quantum chemical
investigations.
a) b) c)
Figure 5.1: Optimized structures ((RI-)BP86/SV(P)) of the focused WCAs: a) = [((CF3)3CO)2B(C6F5)2]−, b) = [((CF3)2HCO)2B(C6F5)2]−, c) = [(CF3CH2O)2B(C6F5)2]−.
104
5.2 Results and Discussion
5.2.1 Quantum Chemical Calculations
We started our studies to design novel borate WCAs [(RFO)2B(C6F5)2]−
(RF = C(CF3)3, CH(CF3)2 and CH2CF3) by performing quantum chemical
calculations to gain an overview over the stability values. Analogously to Chapter
3, we calculated the FIA (fluoride ion affinity), PD (proton decomposition), CuD
(copper decomposition), HOMO level and the HOMO/LUMO gap. Additionally we
performed calculations to obtain ligand affinity values, but in contrast to the
original method, which uses MP2/TZVPP level for calculating the non-isodesmic
reference reaction of the small molecules,[16] we selected the superior g3mp2
method for improved accuracy of typically 5-6 kJ mol−1.[17]
M(L)n− + AlF3 M(L)n-1 + F3Al-L−BP86/SV(P)
g3mp2
ΔH = LA
F3Al-L− AlF3 + L−
M(L)n− M(L)n-1 + L−
Equation 5.1: Definition of the ligand affinity.
Contrary to the homoleptic WCAs highlighted in Chapter 3, the anions
investigated herein are heteroleptic, and thus different decomposition pathways
are possible. Scheme 5.1 shows the approach used to determine the stability
values of a heteroleptic WCA.
[RA2BRB
2]−LAA LAB-RA -RBRABRB
2 RA2BRB
FIAB
[RABFRB2]− [RA
2BFRB]−
FIAA
+Cu+/H+
CuDA/PDA CuDB/PDB
Cu/H-RA Cu/H-RB+ +RABRB2 RA
2BRB
+F− +F−
Scheme 5.1: Procedure to determine the stability values (bold) of a given borate based WCA [RA2BRB
2]−.
105
To simplify matters we defined RB to be -C6F5, since the ligand -C6F5 is part in all
three investigated WCAs. Accordingly, the respective alcoholate residues of the
WCA were labeled as RA. The LAA, CuDA and PDA values reflect the abstraction
of a RA ligand of the WCA and consequently the LAB, CuDB and PDB values
reflect the abstraction of a RB ligand of the WCA. The FIAA value represents the
stability of the Lewis acid that results upon removing the ligand RA from the WCA
and vice versa for the FIAB value.
Table 5.1: Calculated properties of the investigated WCAs [((CF3)3CO)2BRB2]− 8, [((CF3)2HCO)2BRB
2]− 9 and [(F3CH2CO)2BRB
2]− 10, as well of reference anions. RB = C6F5.
Stability value 8 9 10 [Al(OC(CF3)3)4]− [B(C6F5)4]−
FIAA [kJ/mol] 459 435 410 543 452
FIAB [kJ/mol] 455 419 381
PDA [kJ/mol] −1219 −1197 −1211 −1077 −1263
PDB [kJ/mol] −1316 −1331 −1347
CuDA [kJ/mol] −555 −517 −523 −413 −567
CuDB [kJ/mol] −618 −633 −649
LAA [kJ/mol] 220 283 334 342[a] 296[a]
LAB [kJ/mol] 246 222 210
HOMO [eV] −3.152 −3.107 −2.773 −4.096 −3.120
Gap [eV] 4.368 4.303 4.166 6.737 4.214
[a] taken from literature.[16]
Like every compound in nature, a WCA typically decomposes the pathway, which
is energetically favored, so for the stability discussion the adverse values are the
most meaningful.
The higher the FIA value the more stable is a WCA towards ligand abstraction.
The FIAA values herein are always higher than the FIAB values, since the C6F5
group is a potent electron withdrawing group,[18] which enhance the Lewis acidity
of the RFOB(C6F5)2 Lewis acids without being capable of back-bonding.
Additionally upon complexation of the fluoride ion the geometry of the Lewis acid
changes from (pseudo)-trigonal planar to (pseudo)-tetrahedral and thus the
π-B-O stabilization energy of the B-ORF bond gets lost, which is more
pronounced in the Lewis acids (RFO)2BC6F5 with its two B-O bonds. A general
trend emerging from Table 5.1 is the decrease of the FIA values proceeding from
106
left towards right. This trend is based on the continuous substitution of CF3
groups by hydrogen atoms. Although the Lewis acid centre gets more accessible,
since the hydrogen atoms are much less bulky than the CF3 groups, this steric
effect is overcompensated by the loss of the electron withdrawing capability of
the CF3 groups.
The FIAB of compound 8 is similar to the one of [B(C6F5)4]− and other borate
based WCAs (see Table 3.5). Compound 9 represents the structural hybrid
between [B(C6F5)4]− (FIA = 452) and [B(OCH(CF3)2)4]− (FIA = 384), which is
reflected by its FIAB value of 419, lying nearly exact midway between the two
comparative values. The FIAB value of compound 10 is quite low, which is
attributed to the missing electron withdrawing power of the two protons in the
CH2CF3 residues.
The more negative the PD or CuD values are, the more vulnerable is a WCA
against electrophilic attack. The PDB and CuDB values in Table 5.1 are lower than
the respective PDA and CuDA values, which is also based on the stability of the
resulting Lewis acids. As mentioned before, the Lewis acids (RFO)2BC6F5, which
result by an electrophilic attack at the C6F5 residue, are more stable and this is
also reflected in their low PDB and CuDB values. While the PDA and CuDA values
signify a decent stability, the PDB and CuDB values are at the limit of viable
values and thereby the WCAs may be prone to decomposition by electrophiles.
A further tool to rate the stability of a WCA is the ligand affinity (LA), which
represents the energy that is needed to abstract an anionic ligand from the WCA.
The higher the LA value, the more stable is the WCA towards ligand abstraction.
However the LA value depends also on the stability of the resulting anion, which
means if the L− anion is stable the LA value is relatively low.[16] As you can see
the LAA values increase going from left to right in Table 5.1, since the abstracted
alcoholate ligand ORF− becomes less stable. Additionally, the B-ORF bond
strength increases, because less electron density is deducted by the fewer CF3
groups. In contrast the LAB values slightly decrease in the same direction, which
is based on the stability of the generated Lewis acid. Interestingly the critical LA
values (LAA for compound 8, LAB for compound 9 and 10) of the WCAs are in a
similar range.
107
The HOMO energy is used to estimate the resistance versus oxidation and the
HOMO/LUMO gap versus reduction. If the HOMO energy is low, the electron is
harder to remove and to oxidize. A large gap is a measure for the sensitivity of an
anion to reduction and predicts the energy, which is necessary for an electron to
reach the excited state. The HOMO levels and the HOMO/LUMO gap of the
investigated WCAs are in a suitable range.
Summarized it can be said, that the three detailed investigated WCAs own some
theoretically weak points, but in general the stability values are decent and
therefore the synthesis was attempted.
5.2.2 Synthesis and Characterization
The basic starting compound {[Mg2(Et2O)3Br2.4Cl0.6][H2B(C6F5)2]}2 6 for the WCA
synthesis was prepared according to the literature of Wagner et al.[15] This
compound was reacted with two equivalents of the alcohols HOC(CF3)3,
HOCH(CF3)2 and HOCH2CF3 to gain the desired anions.
Preliminary investigations: First we tested the generally ability of the B-H
bonds of compound 6 to react with an alcohol and if both B-H bonds would do so.
Therefore, we added a one percent solution of ethanol in CHCl3 to the unpurified
reaction mixture of 6 and gas formation was observed. After storing the solution
for three days at 2°C we obtained crystals, which were characterized by X-ray
spectroscopy.
+ EtOH {[Mg(OEt2)Cl0.57Br0.43][(EtO)2B(C6F5)2]}2 7Compound 6
The obtained structure is quite similar to the structure of
{[Mg(Et2O)Br][(CH3CH2O)2B(C6F5)2]}2, which was documented in literature by
Serwatowski et al. However, they generated the salt via a different synthesis
route, since they wanted to establish a novel synthesis of the [B(C6F5)4]− anion:[19]
2 B(OEt)3+4 C6F5MgBr {[Mg(OEt2)Br][(EtO)2B(C6F5)2]}2
108
O2
O3
Cl1
O1 B1
Mg1
C6
C8
C7
C5
C2C1
C14C13
C11
C9
C10
C12
C15
C17
C19
C18
C16
C20 F8
F7
F5
C3
F6
F4
F2
F9
F3
F1
F10
Mg1ACl1A
Br1A
Br1
C4
Figure 5.2: Molecular structure of compound 7; thermal ellipsoids are shown at the 50 % probability level H atoms are omitted for clarity. Selected atom distances [pm] and angles [°]: Mg1-Cl1 = 237.6(6), Mg1-Cl1A = 249.5(6), Mg1-Br1 = 264.8(3), Mg1-Br1A = 252.5(3), Mg1-O1 = 205.5(1), Mg1-O2 = 198.3(1), Mg1-O3 = 201.7(1), B1-O1 = 148.6(2), B1-O2 = 149.3(2), B1-C9 = 165.1(2), B1-C15 = 164.5(3); Cl1-Mg1-Cl1A = 86.0(1), Br-Mg1-Br1A = 93.0(1), Mg1-Cl1-Mg1A = 94.0(2), Mg1-Br1-Mg1A = 87.1(1), Mg1-O1-B1 = 96.3(1), Mg1-O2-B1 = 99.2(1), O1-Mg1-O2 = 67.0(1), O1-B1-O2 = 96.9(1), C9-B1-C15 = 115.1(1), O1-B1-C9 = 107.4(1), O1-B1-C15 = 112.7(1), O2-B1-C9 =114.7(1), O2-B1-C15 = 108.6(1).
Figure 5.2 shows the crystal structure of compound 7 (monoclinic, P21/c) together
with selected structural data. In contrast to the before mentioned structure of
{[Mg(Et2O)Br][(EtO)2B(C6F5)2]}2, chlorine statistically occupies the position of
bromine in the (MgX)2 ring. Excluding this ring, the structural data are equal to
thoseof {[Mg(Et2O)Br][(EtO)2B(C6F5)2]}2 within the experimental error.
This promising preliminary investigation of the reactivity paired with the suitable
quantum chemical results were the start for a series of experiments:
109
[H2B(C6F5)2]−
[((CF3)3CO)2B(C6F5)2]− 8 [(F3CH2CO)2B(C6F5)2]− 10[((CF3)2HCO)2B(C6F5)2]− 9
HOCF3
CF3CF3
HOCF3
HCF3
HOCF3
HH
Scheme 5.2: WCA synthesis route using different sterically demanding alcohols.
5.2.2.1 Attempts to synthesize [((CF3)3CO)2B(C6F5)2]− 8
We investigated several different reactions routes to synthesize the desired
[((CF3)3CO)2B(C6F5)2]− WCA, varying the solvent, the number of equivalents of
alcohol and the reaction conditions. It has to be mentioned that most of the
samples still contained quite a lot solid components, which could not be detected
by liquid NMR spectroscopy. The reactions will be discussed in the following.
Reaction in ortho-difluorobenzene (8a-c): To compound 6, which is not
completely soluble in ortho-difluorobenzene (o-DFB), 3.3 equivalents HOC(CF3)3
were added. The sample was kept at 92°C for six hours and NMR
spectroscopically analyzed (reaction 8a, black). Subsequently the reaction
mixture was heated to 92°C for another six hours and again NMR
spectroscopically analyzed (reaction 8b, red). Finally, all volatile components
were removed in vacuo and afterwards the residue was redissolved in Et2O and
CH2Cl2 and analyzed by NMR spectroscopy (reaction 8b, blue).
In the 19F NMR spectrum there are three intensive resonances of perfluoro-tert-
butoxy groups and signals of two different C6F5 ligands. The 19F,19F-COSY of
reaction 8a shows a cross peak of the resonance of a perfluoro-tert-butoxy group
at −72.3 ppm to the signal of the ortho fluorine atoms of one of the C6F5 groups at
−131.2 ppm (Figure 5.3). The splitting pattern of the signal at −72.3 ppm is a
1:4:6:4:1 quintet with a coupling constant of 4.4 Hz indicating the coupling to four
chemically equivalent ortho fluorine atoms (Figure 5.5). This is supported by the
ratio of the integrals of these two signals of 9:4. Consequently, these signals
110
must belong to a molecule with one perfluoro-tert-butoxy- and two C6F5-ligands
attached to the boron atom. The comparably large absolute value of 4.4 Hz of the
formal 7JFF coupling between the CF3 group and the ortho fluorine atoms
indicates that there are considerable contributions of through space interactions
to this coupling constant. Through space spin-spin coupling demands a distance
of the corresponding fluorine atoms that is below the sum of the van der Waals’
radii of two fluorine atoms. The other two resonances of perfluoro-tert-butoxy
groups in the 19F NMR spectrum are singlets.
ppm
-60 -80 -100 -120 -140 -160 ppm
-60
-70
-80
-90
-100
-110
-120
-130
-140
-150
-160
ppm
-70 -71 -72 -73 -74 -75 -76 -77 ppm
-122
-124
-126
-128
-130
-132
-134
-136
-138
Figure 5.3: 19F,19F-COSY correlation (left) and an enlarged section (right) (376.54 MHz) of reaction 8a in o-DFB and d8-toluene at 298 K.
The 11B NMR spectrum (Figure 5.4 below) shows two broad boron resonances at
about 44 and 27 ppm with Δν1/2 of 380 and 1290 Hz respectively. The chemical
shifts as well as the line widths of these signals show that there are no borate
anions present, but only tricoordinate Lewis acids. This means, the mixed
compound that was formed is (CF3)3COB(C6F5)2. After heating the spectra, this
does not change much (Figure 5.4 middle) but when a donor solvent (Et2O) is
added there is a new comparably sharp signal at 2 ppm, which appears at the
shift where the signal of HB(C6F5)2•OEt2 is found later (see below, Figure 5.6).
Therefore it seems reasonable to assign the resonances of the second C6F5
group in the 19F NMR spectra of reaction 8a to HB(C6F5)2, because it must have
already been present before. In the solvent o-DFB this compound probably exists
in its dimeric form.
111
-50-40-30-20-1050 40 30 20 10 0 ppm
Figure 5.4: 11B NMR spectra (128.39 MHz) of reaction 8a (black) and 8b (red) in o-DFB and d8-toluene, as well as approach 8c (blue) in Et2O, CH2Cl2 and CD2Cl2 at 298 K.
After heating (reaction 8b) the intensity of the resonance of the free alcohol at
−74.9 ppm in the 19F NMR spectrum (Figure 5.5 middle) decreases relative to the
intensity of the 19F signals of the other perfluoro-tert-butoxy groups (Figure 5.5
below). Especially the intensity of a second singlet resonance of a perfluoro-tert-
butoxy group at −76.7 ppm increases, compared to the intensity of the resonance
of the free alcohol. This signal belongs probably to an alcoholate function, which
could be attached to Mg cations in the solution. The broadened signals of two
different Et2O molecules found in the 1H NMR spectrum support the presence of
Mg cations in solution. After dissolving in Et2O/CH2Cl2 there are two singlets of
perfluoro-tert-butoxy groups at −76.9 and −77.2 ppm respectively, which are
assigned to alcoholate functions. The quintet of (CF3)3COB(C6F5)2 disappears
and the area of the aromatic resonances show further signals, caused by
hydrolysis.
112
-70 -80 -90 -100 -110 -120 -130 -140 -150 -160 -170 ppm
-72.3 ppm
-77.0 ppm
Figure 5.5: 19F NMR spectra (376.54 MHz) of reaction 8a (black) and 8b (red) in o-DFB and d8- toluene, as well as reaction 8c (blue) in Et2O, CH2Cl2 and CD2Cl2 at 298 K.
Reaction in Et2O/CD2Cl2 (8d): Compound 6 is completely soluble in a mixture of
Et2O and CD2Cl2, in contrast to using only one of these solvents (see Chapter
5.4.2.)
In the 11B NMR spectrum two doublet resonances are visible at 0.0 ppm and −7.4
ppm with a 1JHB coupling constant of 105 Hz and 115 Hz respectively (Figure 5.6
black). In the decoupled 11B{1H} NMR spectrum (Figure 5.6 red) these doublet
splittings vanish, which indicates the presence of only one B-H bond per
molecule. In the 1H,11B-HSQC spectrum (Figure 5.7) the main signal of the 11B
NMR spectrum at 0.0 ppm shows a cross peak to a proton signal at 4.14 ppm
and the resonance at −7.4 ppm a cross peak to the proton signal at 4.00 ppm.
Residues of compound 6 are identified by its 1H NMR signal at 2.18 ppm.
113
-50-40-30-20-1050 40 30 20 10 0 ppm
-50 ppm
Figure 5.6: 11B NMR (black) and 11B{1H} NMR (red) spectra (128.39 MHz) of reaction 8d after heating for 6 h, in Et2O and CD2Cl2 at 298K.
ppm
3.94.04.14.24.34.44.5 ppm
-10
-8
-6
-4
-2
8
6
4
2
0
ppm
1234567 ppm
-15
-10
-5
15
10
5
0
Figure 5.7: 1H,11B-HSQC correlation (left) and an enlarged section (right) (400.17, 128.39 MHz) of reaction 8d after heating for 6 h, in Et2O and CD2Cl2 at 298K.
The 19F NMR spectrum (Figure 5.8) shows signals between −72 and −78 ppm,
belonging to perfluoro-tert-butoxy groups, as well between −132 and −170 ppm,
belonging to fluorine atoms of C6F5 groups of a main product (−134.3, −158.5 and
−165.1 ppm) and a minor product (−135.5, −165.5 and −168.4 ppm). The singlet
resonance at −75.0 is the dominant signal and assigned to the free alcohol.
Furthermore there are two broad resonances at −76.9 and −77.2 ppm, which
were assigned in the NMR discussion of reaction 8b to two different alcoholate
functions. Furthermore there is a singlet signal at −76.1, which does not show up
in samples with the identically composition and therefore is not further explained.
114
In contrast to reaction 8a and 8b (that show clear quintet splitting pattern) there is
a resonance with a doublet of a quintet splitting pattern at −72.3 ppm. The quintet
splitting indicates the coupling to four chemically equivalent ortho fluorine atoms
of two C6F5 groups. The additional splitting is caused by a single proton in the
same molecule.
-60 -70 -80 -90 -100 -110 -120 -130 -140 -150 -160 -170 ppm
-72.3 ppm -160 -162 -164 -166 ppm-134 ppm
Figure 5.8: 19F NMR spectrum (376.54 MHz) of reaction 8d after heating for 6 h in Et2O and CD2Cl2 at 298 K.
Furthermore the 19F,19F-NOESY NMR spectrum (Figure 5.9) shows cross peaks
of this signal at −72.3 ppm with the ortho fluorine atoms of the minor species at
−135.5 ppm. That means, that minor product herein is a [((CF3)3CO)HB(C6F5)2]−
anion. The ratio of the integrals of the perfluoro-tert-butoxy group at −72.3 ppm
and the ortho fluorine atoms at −135.5 ppm is 9:4. Consequently, the minor signal
at −7.4 ppm in the 11B NMR spectrum is assigned to that anion.
115
ppm
-70 -80 -90 -100 -110 -120 -130 ppm-70
-80
-90
-100
-110
-120
-130
Figure 5.9: 19F,19F-NOESY NMR spectrum (376.54 MHz) of reaction 8d after heating for 6 h in Et2O and CD2Cl2 at 298 K. Exchange cross peaks in circles.
The signals of the C6F5 group of the main product show no cross peak in the 19F,19F-NOESY NMR spectrum, but as already mentioned before, the compound
possess one B-H bond. Furthermore, the signal at 4.14 ppm in the 1H NMR
spectrum shows an uneven splitting pattern, which indicates an even number of
chemically equivalent coupling partners. Since the coordinating solvent Et2O is
present, the main compound is HB(C6F5)2•OEt2.
The reactions 8a-c show that the parent Lewis acid (CF3)3COB(C6F5)2 is formed
but not the desired anion. When the starting compounds are heated in
Et2O/CH2Cl2 (reaction 8d), one perfluoro-tert-butoxy group is fixed to give the
[((CF3)3CO)HB(C6F5)2]− anion, but one B-H bond is still remaining untouched
even in the presence of free alcohol. These results show, that the perfluoro-tert-
butoxy group and the C6F5 rings are sterically too demanding to form the
[((CF3)3CO)2B(C6F5)2]− anion.
Further attempts in Et2O/CHCl3 with an excess of alcohol or in toluene were
performed as well, but compound 8 could not be generated.
116
5.2.2.2 Synthesis of [((CF3)2HCO)2B(C6F5)2]− 9
Since the perfluoro-tert-butoxy group turned out as to bulky, we went on to the
sterically less demanding 1,1,1,3,3,3-hexafluoro-2-propoxy group and performed
reactions in Et2O/CD2Cl2 and Et2O.
Compound 6 and 2.2 equivalents HOCH(CF3)2 were completely dissolved in
Et2O/CD2Cl2 and NMR spectroscopically analyzed.
The 11B NMR spectrum (Figure 5.10) shows a narrow signal at 4.1 ppm. The
chemical shift and the half with of 30 Hz indicate that this signal represents a
borate anion. Furthermore there are two doublet resonances visible at 0.0 and
−3.8 ppm with a 1JBH coupling constant of 105 Hz and 100 Hz respectively. The
triplet signal at −30.2 ppm can be assigned to compound 6.
-40-30-20-1040 30 20 10 0 ppm
-20-100 ppm
Figure 5.10: 11B NMR (black) and 11B{1H} NMR (red) spectra (128.39 MHz) of reaction 9 in Et2O and CD2Cl2 at 298K.
Via the cross peaks of the 1H,11B-HSQC NMR spectrum (Figure 5.11) the boron
signals can be related to proton signals. The boron signal at 4.1 ppm correlates
to the septet proton signal (3JHF = 6.4 Hz) at 4.30 ppm, which can be assigned to
the proton of the C-H group of a hexafluoroisopropoxy group and the two signals
117
at 0.0 and −3.8 ppm in the 11B NMR spectrum show cross peaks to proton
signals at 4.14 ppm and 3.81 ppm, which belong to B-H hydrogen atoms.
ppm
-110 9 8 7 6 5 4 3 2 1 0 ppm
-80
-60
-40
-20
80
60
40
20
0
ppm
3.23.43.63.84.04.24.4 ppm
-8
-6
-4
-2
10
8
6
4
2
0
Figure 5.11: 1H,11B-HSQC correlation (left) and an enlarged section (right) (400.17, 128.39 MHz) of reaction 9 in Et2O and CD2Cl2 at 298K.
In the 19F,11B-COSY NMR spectrum (Figure 5.12) all boron signals show cross
peaks to different signals of ortho fluorine atoms of C6F5 groups. The boron signal
at 4.1 ppm can be attached to the signal of the ortho fluorine atoms at −133.3
ppm and the boron signal at 0.0 ppm to the signal of the ortho fluorine atoms at
−134.1 ppm. The boron signal at −3.8 ppm correlates with to the signal of the
ortho fluorine atoms at −134.8 ppm. The compound that gives rise to the latter
resonances has not been identified yet.
118
ppm
-80 -100 -120 -140 -160 ppm
-40
-30
-20
-10
10
0
ppm
-126 -128 -130 -132 -134 -136 -138 -140 ppm
-8
-6
-4
-2
10
8
6
4
2
0
Figure 5.12: 19F,11B-COSY correlation (left) and an enlarged section (right) (376.54 MHz) of reaction 9 in Et2O and CD2Cl2 at 298K.
In a 19F,19F-COSY NMR experiment (appendix, Figure 5.24) the signal of the
ortho fluorine atoms at −133.3 ppm shows a cross peak to the major resonance
at −74.7 ppm, which is in the chemical shift range of CF3 groups. The splitting
pattern (Figure 5.13) is a doublet of a quintet with a coupling constant of 2.8 Hz.
The doublet splitting of 6.4 Hz compares well with the 3JHF coupling constant in a
hexafluoroisopropoxy group and possess the same absolute value as the
coupling found in the septet at 4.30 ppm in the proton NMR spectrum. The
quintet splitting must result from the homonuclear 19F,19F coupling that give rise
to the cross peak in the 19F,19F-COSY spectrum and shows that four equivalent
fluorine atoms couple to the fluorine atoms of the CF3 groups. Consequently, the
corresponding molecule must contain two equivalent C6F5 rings. The CF3 groups
must belong to hexafluoroisopropoxy ligands in the same molecule. The ratio of
the integrals of the signal of the CF3 groups and the resonance of the ortho
fluorine atoms is 6:2, which reveals that the number of hexafluoroisopropoxy
groups and C6F5 groups at this molecule is equal. Without any doubt the
discussed signals represent the target anion [((CF3)2HCO)2B(C6F5)2]− 9 that is
one of two major products in this reaction mixture. The second major product is
easily identified by its already assigned signals as HB(C6F5)2•OEt2 (see Chapter
5.2.2.1).
119
-70 -80 -90 -100 -110 -120 -130 -140 -150 -160 ppm
-74.65 ppm -133.5 ppm
Figure 5.13: 19F NMR spectrum (376.54 MHz) of reaction 9 in Et2O and CD2Cl2 at 298K.
In a further reaction with an excess of HOCH(CF3)2 alcohol in pure Et2O no educt
was NMR spectroscopically detected nor any compounds with a hydrogen atom
directly bond to the boron either. The reaction was obviously complete. Apart
from unreacted alcohol there were only two minor side products containing boron
and C6F5 ligands but no B-H bonds. Since no crystals were obtained, the crude
product could not be purified. However, in this second reaction the molar ratio of
anion 9 in the reaction mixture was enhanced from 39% (obtained in the reaction
in CD2Cl2/Et2O) to 59% (derived from fluorine integrals).
5.2.2.3 Synthesis of [(CF3CH2O)2B(C6F5)2]− 10
In this row the sterically least demanding HOCH2CF3 alcohol was reacted with
compound 6 in EtO2/CH2Cl2. Spontaneous gas formation occurred and stirring
was kept up for 24 h.
NMR Spectroscopic Characterization: The 11B NMR spectrum (Figure 5.14)
shows a main signal at 5.6 ppm with Δν1/2 of 215 Hz. The cross peak in the 1H,11B-COSY NMR spectrum (Figure 5.15) allows to assign the boron signal to
the quartet (1JHF = 9.0 Hz) at 4.10 ppm of a CH2 group of an OCH2CF3 ligand.
120
There is at least a second quartet of another OCH2CF3 group but significantly
smaller intensity.
-60-50-40-30-20-1070 60 50 40 30 20 10 0 ppm
Figure 5.14: 11B NMR spectrum (128.39 MHz) of reaction 10 in Et2O and CD2Cl2 at 298K.
ppm
8 7 6 5 4 3 2 1 0 ppm
-25
-20
-15
-10
-5
25
20
15
10
5
0
ppm
3.23.43.63.84.04.24.44.64.85.0 ppm10
9
8
7
6
5
4
3
2
1
Figure 5.15: 1H,11B-COSY correlation (left; artefact from t1 noise in a circle) and an enlarged section (right) (400.17, 128.39 MHz) of reaction 10 in Et2O and CD2Cl2 at 298K.
The most intensive signal in the 19F NMR spectrum (Figure 5.17) is the
resonance of a CF3 group at −73.4 ppm. This must be the resonance of the
OCH2CF3 group that is attached to the boron atom. The second much less
intensive triplet at −77.0 ppm can only belong to the second OCH2CF3 species
already detected in the proton NMR spectrum. In the 19F,11B-COSY NMR
spectrum (Figure 5.16) the 11B signal at 5.6 ppm shows additionally a cross peak
to the signal of ortho fluorine atoms at −134.3 ppm. The ratio of the integrals of
the signal of the CF3 group of the trifluoroethoxy group at −73.4 ppm and the
121
resonance of the ortho fluorine atoms is 3:2, and so the number of the two
different ligands at this molecule is equal. These findings confirm that the anion
[(CF3CH2O)2B(C6F5)2]− 10 was successfully synthesized. In the 11B NMR
spectrum the signal of 10 is the only one that can be detected, demonstrating that
the reaction was complete. The second signal in the 19F NMR spectrum results
from remaining alcohol.
ppm
-60 -80 -100 -120 -140 -160 ppm
-40
-30
-20
-10
40
30
20
10
0
ppm
-126 -128 -130 -132 -134 -136 -138 ppm
-5
20
15
10
5
0
Figure 5.16: 19F,11B-COSY correlation (left) and an enlarged section (right) (376.54 MHz) of reaction 10 in Et2O and CD2Cl2 at 298K.
-70 -80 -90 -100 -110 -120 -130 -140 -150 -160 ppm
-73.5 ppm
Figure 5.17: 19F NMR spectrum (376.54 MHz) of reaction 10 in Et2O and CD2Cl2 at 298K.
Also, vibrational spectra of 10 were recorded and show the expected bands.
122
After storage at 2°C for two weeks anion 10 could be successfully crystallized as
{[Mg(Et2O)Cl][(CF3CH2O)2B(C6F5)2]}2 (yield: 67 %) and analyzed.
X-ray Crystal Structure Characterization: Figure 5.18 shows the dimeric
structure of compound 10 (monoclinic, P21/n). The two [(CF3CH2O)2B(C6F5)2]−
fragments are connected by a nearly square (MgCl)2 ring, with a sum of the bond
angles of 360° (Cl-Mg-Cl angles: 88.22°; Mg-Cl-Mg angels: 91.78°). Each of the
Mg atoms is additionally coordinated by one Et2O molecule on the opposite sides
and each Mg atom is part of a nearly planar four-membered ring (sum of the
angles: 359°) with the two oxygen atoms and one boron atom of a
[(CF3CH2O)2B(C6F5)2]− fragment. The B-O distances amount to 150 and 151 pm
and are enlarged compared to the B-O distances in compound 7 of 147 and 149
ppm. The B-C distances amount to 164 pm and are equal within the experimental
error.
O2
O3
Cl1
O1
Cl1AB1
Mg1
C6
C4
C8C7
F11
C5
C3
C2C1
C14
C13C11
C9C10
C12
C15
C17
C19
C18
C16C20
F8
F7
F5
F6
F4
F2
F9
F3
F1
F16
F15
F13F14
F12
F10
Mg1A
Figure 5.18: Molecular structure of compound 10; thermal ellipsoids are shown at the 50 % probability level. H atoms are omitted for clarity. Selected atom distances [pm] and angles [°]: Mg1-Cl1 = 239.3(6), Mg1-Cl1A = 242.3(6), Mg1-O1 = 206.5(1), Mg1-O2 = 204.7(1), Mg1-O3 = 200.9(1), B1-O1 = 150.2(2), B1-O2 = 150.8(2), B1-C9 = 164.1(2), B1-C15 = 163.9(2); Cl1-Mg1-Cl1A = 88.2(2), Mg1A-Cl1-Mg1 = 91.8(2), Mg1-O1-B1 = 98.5(8) , O1-Mg1-O2 = 66.1(4), O1-B1-O2 = 96.4(1), C9-B1-C15 = 115.7(1), O1-B1-C9 = 107.3(1), O1-B1-C15 = 114.4(1), O2-B1-C9 = 112.6(1), O2-B1-C15 = 109.0(1).
123
5.2.2.4 Metathesis and Alternative Synthesis Routes
Since the Mg containing cation is not helpful for the solubility of the anions, we
tried to get rid of it by metatheses reactions. The reaction mixtures 8a and 8d
were reacted with AgF and activated in an ultra sonic bath for several days but no
conversion occured. Furthermore, compound 10 was dissolved in mixtures of
different ratios of DMC/PC and LiCl or LiBr were added, to obtain crystals of the
alkali salt, but all attempts so far failed (Scheme 5.3 left).
An alternative possible synthesis route based on the Lewis acid RFOB(C6F5)2 (RF
= (CF3)2HC-), which may be reacted with LiORF to give the desired Li salt. To
generate the Lewis acids we tested the Piers’ borane HB(C6F5)2•SMe2 as
reactant, which can be synthesized easily, starting from compound 6 (Scheme
5.3 top lane).[15] Usually the Piers’ borane finds application to form zwitterionic
olefin polymerization catalysts[20] or in hydroboration reactions.[21] Herein, we
wanted to react the single B-H bond with an alcohol or alcoholate function to give
H2 or LiH and the Lewis acid (Scheme 5.3 right).
{[Mg2(OEt2)3Br2Cl][H2B(C6F5)2]}2
{[Mg2(OEt2)3Br2Cl][(RFO)2B(C6F5)2]}2
Ag/Li[(RFO)2B(C6F5)2]
HB(C6F5)2•SMe2Me3SiCl in SMe2
+ 2 HORF
+ LiORF
+ H/Li-ORF
RFOB(C6F5)2•SMe2
+ AgF/LiCl/LiBr
Scheme 5.3: Alternative attempts to synthesize the anion 9.
Adding HOCH(CF3)2 to HB(C6F5)2•SMe2: No significant reaction occurs when
HB(C6F5)2•SMe2 is reacted with HOCH(CF3)2 at room temperature, but if the
reaction mixture (11) is heated to 140°C for several days the Lewis acid
(CF3)2HCOB(C6F5)2•SMe2 11 is formed.
124
The 11B NMR spectrum (Figure 5.19) shows beside the signal of the educt
HB(C6F5)2•SMe2 at −12.0 ppm a broad main signal (Δν1/2 = 418 Hz) at 43.4 ppm.
The chemical shift of the main signal indicates the presence of a Lewis acid.
-50-40-30-20-1050 40 30 20 10 0 ppm
Figure 5.19: 11B NMR (black) and 11B{1H} NMR (red) spectra (128.39 MHz) of reaction 11 in d8- toluene at 298K.
The main signals in the 19F NMR spectrum are a doublet of a quintet resonance
(3JHF = 5.4 Hz, 6JFF = 2.2 Hz) of the CF3 group of an OCH(CF3)2 residue at –74.8
ppm and the resonances of fluorine atoms of a C6F5 group at –131.2 (ortho-F), –
156.3 (para-F) and –159.5 ppm (meta-F). The 19F,19F-COSY experiment
(appendix, Figure 5.25), confirmed by the splitting pattern of the CF3 group in the 19F NMR spectrum (appendix, Figure 5.26), proof that the mentioned fluorine
main signals belong to one molecule. Additionally the ratio of the integrals of the
CF3 group and the ortho fluorine atoms of the C6F5 groups is 6:4. No boron-
fluorine cross peak was detectable in the chosen 2D NMR experiments, but the
only chemically reasonable compound applying to the found signals is the Lewis
acid (CF3)2HCOB(C6F5)2•SMe2. Furthermore a small resonance with an doublet
of an triplet splitting pattern (3JHF = 5.4 Hz, 6JFF = 1.9 Hz) can be found in the 19F
NMR spectrum at –75.3 ppm, indicating the Lewis acid ((CF3)2HCO)2BC6F5. The
formation of this species shows that the higher temperature promotes the release
of C6F5 residues.
Adding LiOCH(CF3)2 to HB(C6F5)2•SMe2: If HB(C6F5)2•SMe2 and LiOCH(CF3)2
are dissolved in Et2O an interesting reaction (12) sets in:
Figure 5.20 shows the 11B NMR spectrum of the reaction of HB(C6F5)2•SMe2 with
LiOCH(CF3)2. There are three main signals in the spectrum. The narrow
resonance at 4.3 ppm (Δν1/2 = 35 Hz) displays no 1JHB coupling and the chemical
shift is in the area of anion 9 (see Chapter 5.2.2.2). At −11.7 ppm a doublet is
125
visible (1JHB = 111 Hz), which belongs to the educt HB(C6F5)2•SMe2. The triplet
signal at −31.1 ppm possess a 1JBH coupling constant of 82 Hz. The splitting
pattern and its chemical shift indicate the presence of the anion [H2B(C6F5)2]−.
Furthermore a minor doublet signal can be found at −2.7 ppm.
-40-30-20-1040 30 20 10 0 ppm
-20-100 ppm-20-100 ppm
Figure 5.20: 11B NMR (black) and 11B{1H} NMR (red) spectra (128.39 MHz) of reaction 12 in Et2O at 298K.
The cross peak between the boron signal at −31.1 ppm and the proton signal at
1.97 ppm in the 1H,11B-HSQC NMR spectrum (Figure 5.21) further indicates the
presence of the anion [H2B(C6F5)2]−. This assignment is confirmed by involving a 1H,19F-COSY experiment ( Figure 5.22), which shows a cross peak between the
proton signal at 1.97 ppm and the signal at −132.5 ppm of the ortho fluorine
atoms of C6F5 groups with a quartet splitting on the proton dimension, due to the
B-H coupling (1JHB = 82 Hz). Furthermore the 1H,11B-HSQC NMR spectrum
shows the cross peak of the educt at −11.7 (11B) and 3.72 ppm (1H). Additionally
it can be seen that the boron signal at 4.3 ppm couples to the septet signal of a
hexafluoroisopropoxy group at 4.30 ppm in the 1H NMR spectrum. In the 1H,19F-
COSY spectrum the same proton septet correlates to the signal of CF3 groups at
−74.1 ppm and to the resonance of ortho fluorine atoms of a C6F5 group at
−132.3 ppm. The ratio of the integrals of the both last mentioned fluorine signals
is 6:2. This confirms the presence of the anion [((CF3)2HCO)2B(C6F5)2]− 9, which
is unambiguously proofed by splitting pattern of the signal at −74.1 ppm. Its
126
doublet splitting of 6.2 Hz is in accordance with the coupling constant observed in
the septet splitting in the correlating proton signal and results therefore form the 3JHF coupling. The remaining quintet splitting of 2.6 Hz belongs necessarily to the
homonuclear 19F,19F-coupling. Hence, the molecule must contain two equivalent
C6F5 groups and due to the integral ratio also two equivalent OCH(CF3)2
substituents.
ppm
-3-2-18 7 6 5 4 3 2 1 0 ppm
-60
-50
-40
-30
-20
-10
30
20
10
0
ppm
3.23.43.63.84.04.24.44.64.85.0 ppm
-15
-10
-5
15
10
5
0
Figure 5.21: 1H,11B-HSQC correlation (left) and an enlarged section (right) (400.17, 128.39 MHz) of reaction 12 in Et2O at 298K.
ppm
-80 -100 -120 -140 -160 ppm
-1
8
7
6
5
4
3
2
1
0
ppm
-70 -80 -90 -100 -110 -120 -130 ppm
3.6
3.8
4.0
4.2
4.4
4.6
4.8
5.0
5.2
5.4
5.6
5.8
6.0
Figure 5.22: 1H,19F-COSYcorrelation (left) and an enlarged section (right) (400.17, 376.54 MHz) of reaction 12 in Et2O at 298K.
In the 19F NMR spectrum (Figure 5.23) besides the broad signal of [OCH(CF3)2]−
at −78.0 ppm there is a small resonance of one further OCH(CF3)2 group at −74.7
127
ppm. Additionally there are resonances of a C6F5 substituent not yet assigned.
The intensity of the resonance of the ortho fluorine atoms of this C6F5 group and
the signal of the unknown OCH(CF3)2 residue is 4:6 and indicates that these
signals belong to a [((CF3)2HCO)HB(C6F5)2]− anion, that would also explain the
small doublet at −2.7 ppm in the 11B NMR spectrum. Interestingly the molar ratio
of the anions [H2B(C6F5)2]− and [((CF3)2HCO)2B(C6F5)2]− in this solution is 1:1.
After seven days at room temperature the resonances of the educt have totally
vanished but the molar ration of the both mentioned anions of 1:1 remains
unchanged.
-80 -90 -100 -110 -120 -130 -140 -150 -160 -170 -180 ppm
-74.1 ppm -132 ppm
Figure 5.23: 19F NMR spectrum (376.54 MHz) of reaction 12 in Et2O and CD2Cl2 at 298K (red: measured four days earlier).
This observation leads to the assumption that an [((CF3)2HCO)HB(C6F5)2]− anion
is formed, which is rearranging to the higher symmetric anions 9 and 6. Even with
a tenfold excess of LiOCH(CF3)2 the same observations are made, so the
[H2B(C6F5)2]− anion is stable in presence of LiOCH(CF3)2. If the reaction mixture
with small LiOCH(CF3)2 excess is heated at 140°C for several days, the total
decomposition of anion 9 and a partial decomposition of anion 6 were observed.
The Lewis acid (CF3)2HCOB(C6F5)2 cannot be generated following this synthesis
route. However, the anion [((CF3)2HCO)2B(C6F5)2]− 9 is formed with a yield of
50%, which is the maximum that is achievable on this way. So far no crystals of
the lithium salt of the anion have been obtained from this solution.
128
5.3 Conclusion
Herein, possible ways to synthesize the anions [(RFO)2B(C6F5)2]− (ORF =
OC(CF3)3, OCH(CF3)2, OCH2CF3) were investigated and their stability values
calculated and discussed by using the FIA (fluoride ion affinity), PD (proton
decomposition), CuD (copper decomposition), LA (ligand affinity) and the
energetic level of the frontier molecule orbitals. During preliminary investigations
crystals of {[Mg(Et2O)Cl0.57Br0.43][(CH3CH2O)2B(C6F5)2]}2 were obtained and
characterized by XRD.
The synthesis of [((CF3)3CO)2B(C6F5)2]− 8 based on compound 6 was not
successful, since the OC(CF3)3 residues turned out as sterically to demanding.
We just managed to detect the Lewis acid (CF3)3COB(C6F5)2 and the anion
[((CF3)3CO)HB(C6F5)2]−. However, due to the Mg containing cation, insoluble
compounds were formed, which could not be further characterized. Based on
compound 6 the anion [((CF3)2HCO)2B(C6F5)2]− 9 was obtained, but the reaction
was not complete and not free of byproducts. In this row, the conversion of
compound 6 with HOCH2CF3 to the anion [(CF3CH2O)2B(C6F5)2]− 10 was the
cleanest reaction without noteworthy impurities and the crystal structure of
{[Mg(Et2O)Cl][(CF3CHO)2B(C6F5)2]}2 was characterized.
To improve the synthesis of [((CF3)2HCO)2B(C6F5)2]− 9 and to get rid of the Mg
containing cation, an alternative synthesis route based on HB(C6F5)2•SMe2 was
developed. Conversion of the Piers borane with HOCH(CF3)2 alcohol demands
heating to 140°C for several days for the Lewis acid (CF3)2HCOB(C6F5)2•SMe2 to
be formed, but even at this rough reaction conditions the conversion was not
complete. The reaction of HB(C6F5)2•SMe2 with LiOCH(CF3)2 does not exchange
the hydride against an alcoholate residue to form the Lewis acid
(CF3)2HCOB(C6F5)2, but the anion [((CF3)2HCO)HB(C6F5)2]− is generated. This
anion quickly rearranges in solution to the anions [((CF3)2HCO)2B(C6F5)2]− 9 and
[H2B(C6F5)2]− and is detectable in the equilibrium in the solution only in low
quantity.
129
5.4 Experimental Section
5.4.1 Theoretical Methods
All geometries were optimized at the (RI-)BP86/SV(P) level[22] using
TURBOMOLE 6.4.[23] Vibrational frequencies were calculated with the AOFORCE
module[24] and all structures represented true minima without imaginary
frequencies. Ligand affinity values were obtained using the g3mp2 method.[17]
The reference values for calculating the FIA, CuD and PD values were taken
from Chapter 3.
5.4.2 Synthesis and Characterization
Techniques and Instruments: Due to the air and moisture sensitivity of most
materials, all manipulations were undertaken with vacuum and Schlenk
techniques as well as in a glove box with an argon atmosphere (H2O and O2 < 1
ppm). The solvents were dried by using conventional drying agents and directly
distilled. NMR spectra were obtained at room temperature on a BRUKER
AVANCE II+ 400 spectrometer. 1H chemical shifts are given with respect to TMS, 19F NMR spectra to fluorotrichloromethane and 11B NMR spectra to the boron-
trifluoride-diethyl-ether-complex. The IR samples were measured on a Bruker
alpha Fourier transform IR spectrometer arranged in an argon-filled glove box
using a diamond ATR unit (200 - 30000 cm–1). The spectra were recorded in the
range of 375 - 4000 cm–1 and analyzed with the OPUS software package. Raman
spectra were obtained on a BRUKER VERTEX 70 spectrometer with a BRUKER
RAM II Raman unit in sealed melting point capillaries. Data collections for X-ray
structure determinations were performed on a BRUKER APEX II Quazar CCD
area-detector diffractometer with Mo-Kα radiation (λ = 0.71073 Å). The single
crystals were mounted in perfluoroether oil on a MiTeGen MicromountTM.
{[Mg2(Et2O)3Br2.4Cl0.6][H2B(C6F5)2]}2, HB(C6F5)2•SMe2 and LiOCH(CF3)2 were
synthesized according to literature.[15, 25]
130
Reactions to synthesize compound 8 based on compound 6: Reaction in
ortho-difluorobenzene and resulted in Et2O/CH2Cl2 8a-8c: To compound 6 (0.364
g, 0.448 mmol, 1 eq.) in 20 ml o-DFB HOC(CF3)3 (0.21 ml, 0.36g, 1.5 mmol, 3.3
eq.) was added. Compound 6 was not completely soluble in o-DFB. The reaction
mixture was heated for 6 h at 92°C and NMR spectroscopically analyzed.
Subsequently the sample was kept for further 6 h at 92°C and was again NMR
spectroscopically analyzed. The solvent was removed in vacuo and the residue
dissolved in a mixture of 7 ml Et2O and 3 ml CH2Cl2 and analyzed by NMR
spectroscopy. Reaction 8a and 8b: 19F NMR (376.54 MHz, d8-toluene): δ = −72.3
(quint., 9F, 7JFF = 4.4 Hz, (CF3)3COB(C6F5)2), −74.9 (s, 9F, HOC(CF3)3), −76.7 (s,
9F, −OC(CF3)3), −131.2 (m, 4F, o-F in (CF3)3COB(C6F5)2), −132.8 (m, 4F, o-F in
HB(C6F5)2), −139.3 (m, 2F, o-DFB)) −146.6 (m, 2F, p-F in (CF3)3COB(C6F5)2),
−146.9 (m, 2F, p-F in HB(C6F5)2), −161.4 (m, 4F, m-F in HB(C6F5)2), −161.6 (m,
4F, m-F in (CF3)3COB(C6F5)2). 11B NMR (128.39 MHz, d8-toluene): δ = 44 (br,
2B, Δν1/2 = 380 Hz, {HB(C6F5)2}2), 27 (br, 1B, Δν1/2 =1290 Hz, (CF3)3COB(C6F5)2).
Reaction 8c: 19F NMR (376.54 MHz, CD2Cl2): δ = −75.0 (s, 9F, HOC(CF3)3),
−76.9 (s, 9F, −OC(CF3)3), −77.2 (s, 9F, −OC(CF3)3), −133.1 (m, 4F, o-F in
HB(C6F5)2•OEt2), −134.5 (m, 4F, o-F in (CF3)3COB(C6F5)2), −156.5 (t, 2F, p-F in
HB(C6F5)2•OEt2), −162.7 (t, 2F, p-F in (CF3)3COB(C6F5)2), −165.3 (m, 4F, m-F in
HB(C6F5)2•OEt2), −162.7 (m, 4F, m-F in (CF3)3COB(C6F5)2). 11B NMR (128.39
MHz, CD2Cl2): δ = 2 (d, 1B, HB(C6F5)2•OEt2), 27 (br, 1B, Δν1/2 = 1290 Hz,
(CF3)3COB(C6F5)2).
Reaction in Et2O/CH2Cl2 8d: Compound 6 (0.647 g, 0.796 mmol, 1 eq.) was
dissolved in Et2O (20 ml) and CH2Cl2 (10 ml) and HOC(CF3)3 (0.27 ml, 0.46 g, 1.9
mmol, 2.4 eq.) was added and the reaction mixture stirred at RT for 12 h and
subsequently refluxed for 6 h. 1H NMR (400.17 MHz, CD2Cl2): δ = 1.01 (t, 6H,
(CH3CH2)2O), 2.18 (q, 2H, [H2B(C6F5)2]−), 3.31 (q, 4H, (CH3CH2)2O), 4.01 (q, 1H,
[((CF3)3CO)HB(C6F5)2]−), 4.14 (q, 1H, HB(C6F5)2•OEt2), 5.20 (t, 1H, CHDCl2),
5.52 (s, 2H, CH2Cl2), 7.42 (s, 1H, CHCl3), 7.97 (s, 1H, HOC(CF3)3). 19F NMR
(376.54 MHz, CD2Cl2): δ = −72.3 (dquint., 9F, [((CF3)3CO)HB(C6F5)2]−), −75.0 (s,
9H, HOC(CF3)3), −76.1 (unidentified signal), −76.9 (s, 9F, −OC(CF3)3), −77.2 (s,
9F, −OC(CF3)3), −134.3 (m, 4F, o-F in HB(C6F5)2•OEt2), −135.5 (m, 4F, o-F in
[((CF3)3CO)HB(C6F5)2]−), −158.5 (t, 2F, p-F in HB(C6F5)2•OEt2), −165.1 (m, 4F,
131
m-F in HB(C6F5)2•OEt2), −165.5 (t, 2F, p-F in [((CF3)3CO)HB(C6F5)2]−), −168.4 (m,
4F, m-F in [((CF3)3CO)HB(C6F5)2]−). 11B NMR (128.39 MHz, CD2Cl2): δ = 0.0 (d,
1B, 1JBH = 105 Hz, HB(C6F5)2•OEt2), −7.4 (d, 1B, 1JBH = 115 Hz,
[((CF3)3CO)HB(C6F5)2]−).
Synthesis of compound 9 based on compound 6: Compound 6 (0.056 g,
0.069 mmol, 1 eq.) and 0.016 ml HOCH(CF3)2 (0.026g, 0.152 mmol, 2.2 eq.)
were completely dissolved in 1 ml CD2Cl2 and 1 ml Et2O. The reaction mixture
was mixed several minutes and subsequently NMR spectroscopically analyzed.
1H NMR (400.17 MHz, CD2Cl2): δ = 1.15 (t, 6H, (CH3CH2)2O), 2.05 (q, 2H,
[H2B(C6F5)2]−), 3.45 (q, 4H, (CH3CH2)2O), 3.82 (unidentified species), 4.14 (q, 1H, 3JHF = 6.4 Hz, HB(C6F5)2•OEt2), 4.30 (m, 1H, [((CF3)2HCO)2B(C6F5)2]−), 5.34 (t,
1H, CDHCl2). 19F NMR (376.54 MHz, CD2Cl2): δ = −74.7 (dquint., 12F, 3JFH = 6.4
Hz, 7JFF = 2.8 Hz [((CF3)2HCO)2B(C6F5)2]−), −133.3 (d, 4F, o-F in
[((CF3))2HCO)2B(C6F5)2]−), −134.1 (m, 4F, o-F in HB(C6F5)2•OEt2), −158.2 (t, 2F,
p-F in HB(C6F5)2•OEt2), −162.4 (t, 2F, p-F in [((CF3)2HCO)2B(C6F5)2]−), −164.8
(m, 4F, m-F in HB(C6F5)2•OEt2) −167.2 (m, 4F, m-F in [((CF3)2HCO)2B(C6F5)2]−). 11B NMR (128.39 MHz, CD2Cl2): δ = 4.1 (s, 1B, [((CF3)2HCO)2B(C6F5)2]−), 0.0 (d,
1B, 1JBH = 105 Hz, HB(C6F5)2•OEt2), −3.8 (unidentified species), −30.2 (t, 1B,
[H2B(C6F5)2]−).
Synthesis of compound 10 based on compound 6: Compound 6 (0.239 g,
0.294 mmol, 1 eq.) was dissolved in 3 ml Et2O and 4 ml CH2Cl2, and 0.052 ml
2,2,2-trifluoroethanol (0.069 g, 0.686 mmol, 2.9 eq.) was added. Gas formation
occurs and the reaction mixture was stirred for eight hours at RT. Subsequently
the reaction mixture was stored at 2°C for two weeks. A few crystals were
removed and analyzed by XRD. The remaining crystals were filtrated and volatile
compounds removed in vacuo (yield: 0.268 g, 67 %). For NMR analysis the
reaction mixture of an independent sample was used. 1H NMR (400.17 MHz,
CD2Cl2): δ = 1.15 (t, 6H, (CH3CH2)2O), 3.48 (q, 4H, (CH3CH2)2O), 4.10 (q, 4H, 3JHF = 9.0 Hz, [(CF3CH2O)2B(C6F5)2]−), 4.60 (s, 2H, H2), 5.34 (t, 1H, CDHCl2),
5.35 (s, 2H, CH2Cl2), 7.43 (s, 1H, CHCl3) 19F NMR (376.54 MHz, CD2Cl2):
δ = −73.4 (broad triplet, 6F, [(CF3CH2O)2B(C6F5)2]−), −73.4 (br, 3F, CF3CH2O−),
−134.3 (br, 4F, o-F in [(CF3CH2O)2B(C6F5)2]−), −156.1 (br, 2F, p-F in
[(CF3CH2O)2B(C6F5)2]−), −163.6 (m, 4F, m-F in [(CF3CH2O)2B(C6F5)2]−). 11B NMR
132
(128.39 MHz, CD2Cl2): δ = 5.6 (s, 1B, Δν1/2 = 215 Hz, [(CF3CH2O)2B(C6F5)2]−). IR:
ν~ = 403 (vw), 453 (vw), 503 (vw), 603 (vw), 676 (m), 762 (w), 782 (w), 810 (w),
836 (w), 898 (w), 927 (w), 946 (w), 974 (s), 998 (w), 1037 (m), 1084 (vs), 1154
(s), 1163 (s), 1185 (m), 1285 (s), 1319 (w), 1393 (w), 1423 (w), 1451 (vs), 1520
(m), 1648 (w), 2989 (vw), 2230 (vw) cm−1. Raman: ν~ = 169, 256, 285, 354, 393,
410, 449, 485, 497, 515, 584, 678, 784, 811, 836, 896, 933, 1000, 1043, 1087,
1135, 1155, 1290, 1334, 1369, 1396, 1455, 1488, 1504, 1648, 2881, 2940, 2977,
2983 cm−1.
Synthesis of compound 11 based on HB(C6F5)2•SMe2: HB(C6F5)2•SMe2 (0.040
g, 0.098 mmol, 1 eq.) and 0.010 ml HOCH(CF3)2 (0.016 g, 0.098 mmol, 1 eq.)
were dissolved in 1 ml toloul d8. The reaction mixture was heated to 140°C for
seven days and subsequently NMR spectroscopically analyzed. 1H NMR (400.17
MHz, d8-toluene): δ = 1.19 (s, 6H, HB(C6F5)2•SMe2), 1.72 (s, 6H,
((CF3)2HCO)2BC6F5•SMe2), 2.09 (m, 1H, C6D5(CHD2)), 3.55 (m, 1H,
HB(C6F5)2•SMe2), 4.51 (s, 2H, H2), 4.65 (sept, 1H, 3JHF = 5.4 Hz,
((CF3)2HCO)2BC6F5•SMe2), 6.97-7.09 (m, 3H, C6HD4(CD3). 19F NMR (376.54
MHz, d8-toluene): δ = −74.8 (dquint., 6F, 3JHF = 5.4 Hz, 7JFF = 2.2 Hz,
(CF3)2HCOB(C6F5)2•SMe2), −75.3 (dt, 12F, , 3JHF = 5.4 Hz, 7JFF = 1.9 Hz,
((CF3)2HCO)2BC6F5•SMe2), −131.2 (d, 4F, o-F in (CF3)2HCOB(C6F5)2•SMe2),
−131.6 (d, 4F, o-F in HB(C6F5)2•SMe2), −144.6 (t, 2F, p-F in HB(C6F5)2•SMe2),
−156.3 (t, 2F, p-F in (CF3)2HCOB(C6F5)2•SMe2), −159.5 (m, 4F, m-F in
(CF3)2HCOB(C6F5)2•SMe2), −163.1 (m, 4F, m-F in HB(C6F5)2•SMe2). 11B NMR
(128.39 MHz, d8-toluene): δ = 43.4 (br, 1B, Δν1/2 = 418 Hz,
(CF3)2HCOB(C6F5)2•SMe2), −12.0 (d, 1JHB = 98 Hz, HB(C6F5)2•SMe2).
Synthesis of compound 9 based on HB(C6F5)2•SMe2: HB(C6F5)2•SMe2 (0.040
g, 0.098 mmol, 1 eq.) and LiOCH(CF3)2 (0.020 g, 0.115 mmol, 1.2 eq.) were
dissolved in 1 ml Et2O. The reaction mixture was mixed several minutes and
NMR spectroscopically analyzed. 1H NMR (400.17 MHz, no lock): δ = 1.14 (t, 6H,
(CH3CH2)2O), 1.97 (q, 2H, 1JHB = 82 Hz, [H2B(C6F5)2]−), 2.05 (s, 6H, S(CH3)2),
3.41 (q, 4H, (CH3CH2)2O), 3.72 (q, 1H, HB(C6F5)2•SMe2), 4.30 (m, 1H,
[((CF3)2HCO)2B(C6F5)2]−), 4.39 (septet, 1H, 3JFH = 6.9 Hz, −OCH(CF3)2). 19F NMR
(376.54 MHz, no lock): δ = −74.1 (dquint., 12F, 3JFH = 6.4 Hz, 7JFF = 2.8 Hz,
[((CF3)2HCO)2B(C6F5)2]−), −74.1 (d, unidentified species), −78.0 (br., 6F,
133
−OCH(CF3)2), −131.9 (d, 4F, o-F in HB(C6F5)2•SMe2), −132.3 (d, 4F, o-F in
[((CF3)2HCO)2B(C6F5)2]−), −132.5 (br, 4F, o-F in [H2B(C6F5)2]−), −158.1 (t, 2F, p-F
in HB(C6F5)2•SMe2), −162.0 (t, 2F, p-F in[((CF3)2HCO)2B(C6F5)2]−), −164.0 (t, 2F,
p-F in [H2B(C6F5)2]−), −164.4 (m, 4F, m-F in HB(C6F5)2•SMe2), −166.9 (m, 4F, m-
F in[((CF3)2HCO)2B(C6F5)2]−), −167.1 (m, 4F, m-F in [H2B(C6F5)2]−). 11B NMR
(128.39 MHz, no lock): δ = 4.3 (br, 1B, Δν1/2 = 35 Hz [((CF3)2HCO)2B(C6F5)2]−),
−2.7 (d, unidentified species), −11.7 (d, 1B, 1JHB = 111 Hz, HB(C6F5)2•SMe2),
−31.7 (t, 1B, 1JHB = 82 Hz, [H2B(C6F5)2]−).
134
5.5 References
[1] S. H. Strauss, Chemical Reviews 1993, 93, 927; C. A. Reed, Accounts of Chemical Research 1998, 31, 133.
[2] I. Krossing, I. Raabe, Angewandte Chemie-International Edition 2004, 43, 2066.
[3] Rosentha.Mr, Journal of Chemical Education 1973, 50, 331; W. H. Hersh, Inorganic Chemistry 1990, 29, 713.
[4] K. Seppelt, Angewandte Chemie-International Edition in English 1993, 32, 1025; M. Bochmann, Angewandte Chemie-International Edition in English 1992, 31, 1181.
[5] J. H. Golden, P. F. Mutolo, E. B. Lobkovsky, F. J. Disalvo, Inorganic Chemistry 1994, 33, 5374; K. Fujiki, S. Ikeda, H. Kobayashi, A. Mori, A. Nagira, J. Nie, T. Sonoda, Y. Yagupolskii, Chemistry Letters 2000, 62.
[6] A. G. Massey, A. J. Park, Journal of Organometallic Chemistry 1964, 2, 245.
[7] H. Nishida, N. Takada, M. Yoshimura, T. Sonoda, H. Kobayashi, Bulletin of the Chemical Society of Japan 1984, 57, 2600.
[8] W. E. Geiger, F. Barriere, Accounts of Chemical Research 2010, 43, 1030; H. J. Gericke, N. I. Barnard, E. Erasmus, J. C. Swarts, M. J. Cook, M. A. S. Aquino, Inorganica Chimica Acta 2010, 363, 2222.
[9] W. Beck, K. Sünkel, Chemical Reviews 1988, 88, 1405. [10] M. J. Taylor, A. J. Saunders, M. P. Coles, J. R. Fulton, Organometallics,
30, 1334. [11] J. M. Farrell, J. A. Hatnean, D. W. Stephan, Journal of the American
Chemical Society 2012, 134, 15728. [12] A. Bernsdorf, H. Brand, R. Hellmann, M. Kockerling, A. Schulz, A.
Villinger, K. Voss, Journal of the American Chemical Society 2009, 131, 8958.
[13] L. Jia, X. M. Yang, C. Stern, T. J. Marks, Organometallics 1994, 13, 3755; V. C. Williams, G. J. Irvine, W. E. Piers, Z. M. Li, S. Collins, W. Clegg, M. R. J. Elsegood, T. B. Marder, Organometallics 2000, 19, 1619.
[14] R. E. LaPointe, G. R. Roof, K. A. Abboud, J. Klosin, Journal of the American Chemical Society 2000, 122, 9560.
[15] A. Schnurr, K. Samigullin, J. M. Breunig, M. Bolte, H. W. Lerner, M. Wagner, Organometallics 2011, 30, 2838.
[16] I. Krossing, I. Raabe, Chemistry-a European Journal 2004, 10, 5017. [17] L. A. Curtiss, P. C. Redfern, K. Raghavachari, V. Rassolov, J. A. Pople,
Journal of Chemical Physics 1999, 110, 4703. [18] M. P. Thornberry, N. T. Reynolds, P. A. Deck, F. R. Fronczek, A. L.
Rheingold, L. M. Liable-Sands, Organometallics 2004, 23, 1333. [19] R. Anulewicz-Ostrowska, T. Klis, J. Serwatowski, Main Group Metal
Chemistry 2002, 25, 501. [20] Y. M. Sun, R. Spence, W. E. Piers, M. Parvez, G. P. A. Yap, Journal of the
American Chemical Society 1997, 119, 5132; K. D. Conroy, P. G. Hayes, W. E. Piers, M. Parvez, Organometallics 2007, 26, 4464.
[21] J. G. Yu, G. Kehr, C. G. Daniliuc, G. Erker, European Journal of Inorganic Chemistry 2013, 3312; A. Stute, G. Kehr, R. Fröhlichz, G. Erker, Chemical Communications 2011, 47, 4288; C. Chen, F. Eweiner, B. Wibbeling, R. Fröhlich, S. Senda, Y. Ohki, K. Tatsumi, S. Grimme, G. Kehr, G. Erker,
135
Chemistry-an Asian Journal 2010, 5, 2199; D. J. Parks, W. E. Piers, Tetrahedron 1998, 54, 15469.
[22] J. P. Perdew, Physical Review B 1986, 33, 8822; J. P. Perdew, Physical Review B 1986, 34, 7406; A. D. Becke, Physical Review A 1988, 38, 3098.
[23] A. Schäfer, H. Horn, R. Ahlrichs, Journal of Chemical Physics 1992, 97, 2571.
[24] P. Deglmann, F. Furche, Journal of Chemical Physics 2002, 117, 9535; P. Deglmann, F. Furche, R. Ahlrichs, Chemical Physics Letters 2002, 362, 511.
[25] A. Reisinger, N. Trapp, I. Krossing, Organometallics 2007, 26, 2096.
136
137
5.6 Appendix
Legend to the following table and calculations herein.
BP86 BP86 density functional
h Hartree, 1H = 2625.5 kJ mol−1
Evrt Sum of vibrational, rotational and translational
energy including ZPE (= FreeH energy)
ZPE Zero point vibrational energy
S Entropy
H Enthalpy
G Gibbs Energy
Table 5.2: Part 1of the summary the energetic data. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values. Gibbs energies (G) were used to calculate the CuD and PD values. RB = C6F5.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
H(g3mp2)
[h]
G(g3mp2)
[h]
[((CF3)3CO)2BRB2]− −3730.520249 672.22 1.14045 −9793843.520255 −9794183.545422
[((CF3)2HCO)2BRB2]− −3056.943188 634.86 1.05167 −8025397.570566 −8025711.125977
[(CF3CH2O)2BRB2]− −2383.328444 593.07 0.89834 −6256857.114834 −6257124.954905
(CF3)3COBRB2 −2604.937085 498.41 0.89566 −6838787,478654 −6839054,519683
((CF3)3CO)2BRB −3003.124976 524.64 0.96952 −7884207.537566 −7884496.599954
[F3AlOC(CF3)3]− −1667.513506 209.94 0.59000 −4377860.965653 −4378036.874153 −1666.882072 −1666.949603
[F3AlRB]− −1269.311498 182.73 0.51814 −3332404.822154 −3332559.305595 −1268.672465 −1268.728081
(CF3)2HCOBRB2 −2268.153716 479.82 0.84225 −5954577.962885 −5954829.079723
((CF3)2HCO)2BRB −2329.557464 488.39 0.85048 −6115785.547559 −6116039.118171
[F3AlOCH(CF3)2]− −1330.706764 190.46 0.53103 −3493590.976202 −3493749.302797 −1330.141666 −1330.203667
CF3CH2COBRB2 −1931.350800 459.45 0.76091 −5070318.910211 −5070545.775527
(CF3CH2CO)2BRB −1655.947192 446.23 0.70990 −4347257.203633 −4347468.860318
[F3AlOCH2CF3]− −993.895559 168.93 0.44864 −2609311.320675 −2609445.082691 −993.399373 −993.451822
[(CF3)3COB(F)RB2]− −2704.863661 503.60 0.91785 −7101140.511895 −7101414.168872
[((CF3)3CO)2B(F)RB]− −3103.049937 529.45 0.98805 −8146556.709820 −8146851.296928
[(CF3)2HCOB(F)RB2]− −2368.071154 484.74 0.86821 −6216907.275266 −6217166.132077
[((CF3)2HCO)2B(F)RB]− −2429.468489 492.51 0.87246 −6378098.823857 −6378358.947806
[CF3CH2OB(F)RB2]− −2031.258389 463.20 0.78782 −5332623.533944 −5332858.422477
[(CF3CH2O)2B(F)RB]− −1755.843401 449.56 0.72651 −4609532.369325 −4609748.978281
CuOC(CF3)3 −2765.997220 178.06 0.51108 −7261972.822385 −7262125.200887
CuOCH(CF3)2 −2429.191808 158.57 0.44697 −6377706.335125 −6377839.599230
CuOCH2CF3 −2092.382933 137.05 0.36816 −5493432.785986 −5493542.552890
CuC6F5 −2367.834518 152.82 0.43032 −6216617.905607 −6216746.205515
HOC(CF3)3 −1126.042986 202.27 0.46845 −2956232.370688 −2956372.039055
138
139
Table 5.3: Part 2 of the summary the energetic data. The thermal energy corrections are calculated at 298.15 K. Thermal-energy-corrected enthalpies (H) were used to calculate the affinity values. Gibbs energies (G) were used to calculate the CuD and PD values.
Compound (RI-)BP86/SV(P) energy
[h]
Evrt
[kJ mol−1]
Entropy S
[kJ mol−1]
Enthalpy H
[kJ mol−1]
Gibbs Energy G
[kJ mol−1]
H(g3mp2)
[h]
G(g3mp2)
[h]
HOCH(CF3)2 −789.243859 182.69 0.40433 −2071982.473971 −2072103.024960
HOCH2CF3 −452.438780 161.96 0.32169 −1187718.101005 −1187814.012879
HC6F5 −727.892696 176.03 0.38854 −1910911.044103 −1911026.887304
AlF3 −541,889828 31.84 0.2792 −1422702.843352 −1422786.098758 −541.527186 −541.560463
[OC(CF3)3]− −1125,497435 166.01 0.4721 −2954836.282395 −2954977.047954 −1125.232079 −1125.232079
[OCH(CF3)2]− −788.6764670 144.18 0.40019 −2070531.291913 −2070650.608562 −788.480684 −788.480684
[OCH2CF3]− −451.8438420 119.92 0.31689 −1186198.126650 −1186292.607403 −451.718103 −451.718103
[C6F5]− −727.2994842 139.89 0.38527 −1909389.699802 −1909504.568053 −727.026525 −727.070242
Figure 5.24: 19F,19F-COSY NMR spectrum (left) and an enlarged section (right) (376.54 MHz) of reaction 9 in CD2Cl2 and Et2O at 298K.
ppm
-80 -100 -120 -140 -160 ppm-70
-80
-90
-100
-110
-120
-130
-140
-150
-160
ppm
-70 -80 -90 -100 -110 -120 -130 ppm-70
-80
-90
-100
-110
-120
-130
Figure 5.25: 19F,19F-COSY NMR spectrum (left) and an enlarged section (right) (376.54 MHz) of reaction 11 in d8-toluene at 298 K.
ppm
-70 -80 -90 -100 -110 -120 -130 ppm-70
-80
-90
-100
-110
-120
-130
ppm
-60 -80 -100 -120 -140 -160 ppm-60
-70
-80
-90
-100
-110
-120
-130
-140
-150
-160
-70 -80 -90 -100 -110 -120 -130 -140 -150 -160 -170 ppm
-74.75 ppm
Figure 5.26: 19F NMR spectrum (376.54 MHz) of reaction 11 in d8-toluene at 298K.
140
6 Conclusion
In the past years Lewis acids, and especially their strongest representatives,
were of great interest and found widespread applications. Currently frustrated
Lewis pair (FLP) chemistry begins to shine, since FLPs are metal-free and
capable to activate small molecules like H2, alkenes, aldehydes and CO2.[1] In this
thesis the possibility to obtain frustrated Lewis pairs based on the Lewis acid
B(Ohfip)3 1 (Ohfip = OCH(CF3)2) was investigated. Before turning to FLP
chemistry experiments, the thermodynamics of the hydrogen activation of 1 in
comparison to the well-known B(C6F5)3 was analyzed by using a Born-Fajans-
Haber cycle. The investigations were continued by experiments, reacting 1 with
selected phosphanes, amines, pyridines and N-heterocyclic carbenes, but no
reaction was observed. Even upon addition of very strong Lewis bases, adduct
formation was in part only occurring in equilibrium. However, the crystal
structures of the very weak adduct 1•NCMe was analyzed, which possesses to
the best of our knowledge the longest coordination between boron and a nitrogen
atom of an acetonitrile molecule (248 pm). Surface analyses were performed,
which underline the weakness of this coordination.
Figure 6.1: Visualized surface analyses (left: Hirshfeld surface; right: electrostatic potential on an electro density surface) of 1•NCMe.
Induced by these unexpected unsuccessful FLP reactions, calculations were
performed to investigate and compare the Lewis acidity towards HSAB hard and
141
soft ions for a deeper insight. Therefore, a unified reference system as anchor
point was chosen. This reference system is based on the trimethylsilyl
compounds Me3SiY (Y = F, Cl, H, Me) of the respective ions, so that the relative
values of the Lewis acidity towards different bases are better comparable. A
further advantage is that only the respective reference reaction has to be
optimized at a sufficiently accurate and highly correlated level (here G3), while
the residual in part very large molecules can be assessed based on subsequent
isodesmic reactions calculated at a much less expensive level (here (RI-
BP86/SV(P)). By using this method to gain Lewis acidity scales, 33 common and
frequently used Lewis acids were ranked with respect to their FIA (fluoride ion
affinity), CIA (chloride ion affinity), HIA (hydride ion affinity) and MIA (methyl ion
affinity). The relationship between the ions affinity values and the LUMO levels in
respect of their chemical hardness were analyzed and discussed as well.
Furthermore, the known scales for weakly coordinating anion (WCA) stability
were expanded with the PD (proton decomposition) and CuD (copper
decomposition) of a series of hitherto not explored WCAs.[2]
A further spotlight in this thesis was on the synthesis and characterization of
ferrocenylboranes. As ligand -Ohfip was chosen. Due to its bulky structure,
paired with electron-withdrawing CF3 groups and an oxygen atom capable to π-
interact with boron, this ligand is predestinated for investigations to deepen the
knowledge of ferrocenylboranes. Herein the straightforward synthesis and
characterization of [1-(BOhfip2)Fc] 2 and [1,1’-(BOhfip2)2)Fc] 3 and the
subsequent oxidation of 2 and 3 with Ag+[Al(ORF)4]− to the ferrocinium derivates
[1-(BOhfip2)Fc]+[Al(ORF)4]− 4 and [1,1’-(BOhfip2)2)Fc]+[Al(ORF)4]− 5 was
performed (Scheme 6.1). Furthermore, the crystal structures of 2 and 3 and even
the rare crystal structures of the oxidized ferrocenylboranes 4 and 5 were
obtained. It could be shown that the dip angle α* (α* = 180° − α, with α being the
angle Cp(centroid)-Cipso-B)) strongly decrease upon oxidation and the Fe-B
interaction is enfeebled.
To further investigate the electron-withdrawing potential of the B(Ohfip)2 group,
cyclic voltammetric experiments were run on the compounds 2 and 3, showing
the expected anodic shifts. Additionally, the CIA, HIA, FIA and MIA values of the
neutral compounds 2 and 3 have been determined accordingly to the previously
142
introduced method. The analyses of the ferrocenylboranes were completed by
NMR-, IR-, RAMAN- and UV/VIS-spectroscopy.
1. BBr3
2. LiOhfip
1. 2 BBr3
2. 2 LiOhfip
[1-(BOhfip2)Fc] 2 [1,1’-(BOhfip2)2)Fc] 3
[1-(BOhfip2)Fc]+ 4 [1,1’-(BOhfip2)2)Fc]+ 5
Ag+[Al(ORF)4]− Ag+[Al(ORF)4]−
Scheme 6.1: Schematic representation of the synthesized compounds 2-5 starting from ferrocene. F atoms, H atoms of the ferrocenyl moieties and [Al(ORF)4]− anions (ORF = O(C(CF3)3) are omitted for clarity. The structure of ferrocene was taken from literature.[3]
Furthermore, possible ways to synthesize the anions [(RFO)2B(C6F5)2]−
(ORF = OC(CF3)3, OCH(CF3)2, OCH2CF3) were investigated and their stability
values calculated and discussed by using the FIA, PD, CuD, LA (ligand affinity)
as well as the energetic level of the frontier molecule orbitals. All of the reactions
were monitored by extensive NMR spectroscopy. During preliminary
investigations crystals of {[Mg(Et2O)Cl0.57Br0.43][(CH3CH2O)2B(C6F5)2]}2 7 were
obtained and characterized by XRD.
It could be shown, that the OC(CF3)3 ligand is sterically too demanding to form
the anion [((CF3)3CO)2B(C6F5)2]− 8. In contrast the anion [((CF3)2HCO)2B(C6F5)2]−
143
9 is accessible via two different synthesis routes, however until now not in very
good yields. Especially the reaction pathway starting from HB(C6F5)2•SMe2
(Scheme 6.2 blue) is considered to be promising for further investigations. The
synthesis of [(CF3CH2O)2B(C6F5)2]− 10 was the cleanest reaction without
noteworthy impurities and crystals suitable for XRD were obtained. However its
stability values are the poorest in that row.
[H2B(C6F5)2]−
[((CF3)3CO)2B(C6F5)2]]− 8
HOC(CF3)3 HOCH(CF3)2 HOCH2CF3
[((CF3)2HCO)2B(C6F5)2]]− 9 (59%)
[(CF3H2CO)2B(C6F5)2]]− 10
HB(C6F5)2•SMe2
[H2B(C6F5)2]− (50%)+
(50%)
LiOCH(CF3)2
(CF3)2HCOB(C6F5)2•SMe2 11
HOCH(CF3)2
Scheme 6.2: Reaction routes carried out to obtain the anions 8-10.
Summarized, the Ohfip ligand has shown interesting characteristics. It has
proven its electron-withdrawing character considering the anodic shift and α*
angle in ferrocenylboranes. In contrast, under the used conditions the Lewis acid
B(Ohfip)3 does not form FLP systems capable to cleave hydrogen and its Lewis
affinity values are quite low. Furthermore, oxidizing the ferrocenylboranes to their
respective ferrocinium derivates reveals, that the contribution to stabilize the
Lewis acid center originates mainly from the π-B-O interaction and not the π-
Cipso-B interaction. Since the former low BO2-/Cp-plane angles in 2 and 3 strongly
increase upon oxidation, leaving the π-B-O geometry and interaction nearly
unchanged, but making a π-Cipso-B interaction impossible. Finally, the decent
144
sterical demand paired with the electronic properties of the Ohfip ligand can be
utilized in fine tuning the composition of ligands in WCAs.
6.1 References
[1] D. W. Stephan, G. Erker, Angewandte Chemie-International Edition 2010, 49, 46.
[2] I. Krossing, I. Raabe, Chemistry-a European Journal 2004, 10, 5017. [3] F. Takusagawa, T. F. Koetzle, Acta Crystallographica Section B-Structural
Science 1979, 35, 1074.
145
7 Crystal Structure Data
Table 7.1: Crystal structure data of the compounds 1, 1•NCMe, and 2.
Compound 1 1•NCMe 2CCDC deposition # 1004582 1004583 865983
Empirical formula C9H3BF18O3 C11H4BF18NO3 C22H12B2F24FeO4
Formula weight 511.92 550.96 873.79
Temperature [K] 120(2) 110(2) 100(2)
Crystal system monoclinic trigonal orthorhombic
Space group P21/c R-3c Pcab
a [Å] 8.923(1) 13.079(2) 8.399(1)
b [Å] 15.103(2) 13.079(2) 17.572(1)
c [Å] 12.973(2) 18.570(2) 25.496(1)
α [°] 90 90 90
β [°] 105.34(3) 90 90
γ [°] 90 120 90
Volume [Å3] 1686.2(4) 2751.0(10) 3763.5(2)
Z 4 6 1
Density [mg mm−3] 2.017 1.995 0.386
μ [mm−1] 0.270 0.257 0.136
F(000) 992 1608 428
θ range for data 2.11 to 24.72 3.12 to 27.45 1.97 to 32.55
collection [°]
Index ranges −10<=h<=10 −16<=h<=16 −10<=h<=12
−17<=k<=17 −16<=k<=16 −26<=k<=26
−15<=l<=15 −24<=l<=24 −38<=l<=38
Reflections collected 13673 14414 49713
Independent reflections 2882 1399 6843
[Rint = 0.0730] [Rint = 0.0372] [Rint = 0.0283]
Data/restraints/parameters 2882/0/280 1399/2/109 6843/0/289
Goodness-of-fit on F2 1.024 1.083 1.036
Final R indexes R1 = 0.0425 R1 = 0.0246 R1 = 0.0317
[I>2σ(I)] wR2 = 0.0808 wR2 = 0.0596 wR2 = 0.0777
Final R indexes R1 = 0.0979 R1 = 0.0262 R1 = 0.0450
[all data] wR2 = 0.0997 wR2 = 0.0607 wR2 = 0.0847
Largest diff. peak/hole 0.24/−0.27 0.17/−0.16 0.50/−0.41
[e Å−3]
146
Table 7.2: Crystal structure data of the compounds 3, 4 and 5.
Compound 3 4 5CCDC deposition # 866159 1004581 865569
Empirical formula C88H48B8F96Fe4O16 C32H11AlBF48FeO6 C38H12AlB2F60FeO8
Formula weight 3495.14 1497.05 1840.93
Temperature [K] 110(2) 110(2) 110(2)
Crystal system monoclinic monoclinic monoclinic
Space group C2/c P21/c P21/c
a [Å] 24.141(2) 21.643(1) 15.381(2)
b [Å] 9.5346(8) 19.973(1) 26.752(2)
c [Å] 12.8764(8) 27.927(2) 15.092 (2)
α [°] 90 90 90
β [°] 94.51(1) 128.65 109.409(2)
γ [°] 90 90 90
Volume [Å3] 2954.6(4) 9428.3(5) 5857.4(5)
Z 4 8 4
Density [mg mm−3] 1.964 2.109 2.088
μ [mm−1] 0.694 0.574 0.514
F(000) 1712 5816 3572
θ range for data 3.17 to 27.48 1.58 to 26.47 1.40 to 22.76
collection [°]
Index ranges −27<=h<=31 −27<=h<=27 −15<=h<=16
−12<=k<=12 −24<=k<=24 −29<=k<=26
−16<=l<=16 −34<=l<=34 −16<=l<=15
Reflections collected 27288 145080 38193
Independent reflections 3385 19320 7881
[Rint = 0.0875] [Rint = 0.1209] [Rint = 0.0491]
Data/restraints/parameters 3385/0/264 19320/0/1603 7881/0/991
Goodness-of-fit on F2 1.075 0.904 1.068
Final R indexes R1 = 0.0425 R1 = 0.0482 R1 = 0.0620
[I>2σ(I)] wR2 = 0.0877 wR2 = 0.0678 wR2 = 0.1529
Final R indexes R1 = 0.0568 R1 = 0.1256 R1 = 0.0900
[all data] wR2 = 0.0942 wR2 = 0.0886 wR2 = 0.1737
Largest diff. peak/hole 0.40/−0.47 0.42/−0.42 1.62/−0.65
[e Å−3]
147
Table 7.3: Crystal structure data of the compounds 7 and 10.
Compound 7 10CCDC deposition # 1004584 1004585
Empirical formula C40H40B2Br0.86Cl1.14F20Mg2O6 C40H28B2C12F32Mg2O6
Formula weight 1176.08 1353.76
Temperature [K] 100(2) 100(2)
Crystal system monoclinic monoclinic
Space group P21/c P21/c
a [Å] 12.600(1) 11.436(2)
b [Å] 7.927 (2) 9.042(1)
c [Å] 23.630(2) 24.565(5)
α [°] 90 90
β [°] 93.48(1) 96.75(2)
γ [°] 90 90
Volume [Å3] 2355.70(10) 2522.52(19)
Z 2 2
Density [mg mm−3] 1.658 1.782
μ [mm−1] 0.966 0.319
F(000) 1183 1344
θ range for data 2.30 to 30.37 1.67 to 29.57
collection [°]
Index ranges −17<=h<=17 −15<=h<=15
−11<=k<=11 −12<=k<=11
−32<=l<=32 −34<=l<=30
Reflections collected 25745 34072
Independent reflections 6633 7073
[Rint = 0.0352] [Rint = 0.0257]
Data/restraints/parameters 6633/536/33 7076/28/397
Goodness-of-fit on F2 1.021 1.018
Final R indexes R1 = 0.0355 R1 = 0.0365
[I>2σ(I)] wR2 = 0.0699 wR2 = 0.0901
Final R indexes R1 = 0.0601 R1 = 0.0435
[all data] wR2 = 0.0766 wR2 = 0.0946
Largest diff. peak/hole 0.36/−0.27 1.07/−0.73
[e Å−3]
148
8 Abstract
This thesis deals with the experimental and theoretical study of Lewis acids and
Weakly Coordinating Anions (WCAs), especially their boron-based
representatives.
In the past years, Lewis acids were of great interest and found widespread
applications. Ranking Lewis acidity is therefore a very useful and powerful tool to
evaluate the abilities of any given Lewis acid. In this work, a novel theoretical
Lewis acid scale based on a unified reference system (Me3SiY; Y = F, Cl, H, Me)
as anchor point was introduced and validated by high level CCSD(T) calculations.
Using this method, the most common Lewis acids were ranked and compared
with respect to the fluoride, chloride, hydride and methyl ion as bases. In this
context, the ability of B(Ohfip)3 (hfip = CH(CF3)2) to activate hydrogen in terms of
FLP (Frustrated Lewis Pair) chemistry was investigated.
Since the Fe-B interaction in ferrocenylboranes is still not well understood and
crystal structures of their oxidized species are rare, the synthesis and
characterization of [1-(BOhfip2)Fc] and [1,1’-(BOhfip2)2)Fc] and the subsequent
oxidation with Ag+[Al(ORF)4]− (RF = C(CF3)3)) to the ferrocinium derivates
[1-(BOhfip2)Fc]+[Al(ORF)4]− and [1,1’-(BOhfip2)2)Fc]+[Al(ORF)4]− are reported
herein. Overall, the obtained compounds enabled a deeper insight into the Fe-B
interaction in ferrocenylboranes, as well as into the electronically and steric
properties of the Ohfip ligand.
In a series of experiments, the ligands (L) OC(CF3)3, OCH(CF3)2 and OCH2CF3
were reacted with different starting materials to give the WCAs [L2B(C6F5)]−. New
knowledge in regarding the synthesis and stability of the objectives could be
gained, thus showing a clear trend based on steric demand.
The presented results are supported by quantum chemical calculations.
Keywords: Lewis acids, Lewis acid scale, boron, ferrocenylboranes, oxidation,
weakly coordinating anions, frustrated Lewis pair.
149
9 Kurzzusammenfassung
Die vorliegende Arbeitet beschäftigt sich mit der praktischen und theoretischen
Untersuchung von Lewis Säuren und schwach koordinierende Anionen (WCAs),
insbesondere mit deren borbasierenden Vertretern.
In den letzten Jahren stieg das Interesse an Lewis Säuren kontinuierlich und
selbige fanden weitgefächerte Anwendungsgebiete. Deshalb ist die Einordnung
von Lewis Säuren in Skalen ein nützliches Werkzeug, um die Eigenschaften
einer Lewis Säure abschätzen zu können. In dieser Arbeit wurde eine neue
theoretische Lewis Säure-Skala eingeführt, die auf einem einheitlichen
Referenzsystem (Me3SiY; Y = F, Cl, H, Me) aufgebaut ist und mit Hilfe von
aufwendigen Berechnungen validiert wurde. Durch diese Methode wurden
weitverbreitete Lewis Säuren hinsichtlich ihrer Fluorid-, Chlorid-, Hydrid- und
Methyl-Anionen Affinität eingeordnet und miteinander vergleichen. In diesem
Zusammenhang wurde untersucht, inwiefern die Lewis Säure B(Ohfip)3 (hfip =
CH(CF3)2) im Sinne der FLP Chemie (Frustrierte Lewis Paare) fähig ist
Wasserstoff zu aktivieren.
Da die Fe-B Wechselwirkung in Ferrocenylboranen noch nicht gut verstanden ist
und noch nicht viele Kristallstrukturen ihrer oxidierten Verbindungen veröffentlicht
wurden, wird hier die Synthese und Charakterisierung von [1-(BOhfip2)Fc] und
[1,1’-(BOhfip2)2)Fc], sowie die darauf folgenden Oxidation mit Ag+[Al(ORF)4]− (RF
= C(CF3)3)) zu den entsprechenden Ferrociniumverbindungen vorgestellt. Die
erhaltenen Verbindungen ermöglichen einen tieferen Einblick in die Fe-B
Wechselwirkung in Ferrocenylboranen, sowie in die elektronischen und
sterischen Eigenschaften des Ohfip Liganden.
In einer ausgedehnten Versuchsreihe wurden die Liganden (L) OC(CF3)3,
OCH(CF3)2 und OCH2CF3 mit verschiedenen Edukten umgesetzt um die
entsprechenden WCAs [L2B(C6F5)]− zu erhalten. Hinsichtlich der Synthese und
Stabilität der Zielverbindungen konnte neues Wissen gewonnen werden.
Die aufgezeigten Ergebnisse werden von begleitenden quantenchemischen
Rechnungen gestützt.
150
151
Stichwörter: Lewis Säuren, Lewis Säure-Skala, Bor, Ferrocenylborane,
Oxidation, schwach koordinierende Anionen, frustrierte Lewis Paare.
