Lewis Acids and Bases

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MSc Chemistry

Molecular Sciences

Literature Thesis

Lewis Acidity

Physical properties, Strength and

Applications

by

Donovan Cerneüs

11942061

May 2019

12 EC

March 2019 – May 2019

Supervisor/Examiner: Examiner:

dhr. dr. J.C. (Chris) Slootweg

prof. dr. B. (Bas) de Bruin

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Abstract

Lewis acidity is the physical organic concept of the ability of a molecule to accept a pair of electrons. The strength of a Lewis acid is dependent on various basic physical organic principles like electrophilicity, inductive effects, sterics and molecular orbital overlap to mention a few. While Brønsted acids have a well-established and widely used scale of acidity strength, better known as the pH scale, Lewis acids do not. Some considerable effort has been put into the development of a Lewis acidic scale, but the complex nature of Lewis acidity makes this exceptionally difficult. The widely accepted fluoride ion affinity (FIA) is a measurement of Lewis acidity, but it has its drawbacks, just as the more recently proposed global electrophilicity index (GEI). The simple fluorescence using method of determining Lewis acidity might be the next step into the right direction.

One feature in particular makes it quite difficult to have an acidity rating of Lewis acidity and that is the affinity of Lewis acids to bases differs from soft- to hard-Lewis acids. Molecular orbital overlap is closely linked with the hard-soft theory and serves that softer species interact better with softer Lewis acids and vice versa. To get a true overview of Lewis acidity it might be best to combine multiple theories taking the hard-soft interactions into account.

Understanding the basic physical organic principles behind Lewis acidity helped scientists in creating extreme Lewis acidities better known as Lewis super acids. These are typically defined as molecular Lewis acids that are stronger than monomeric SbF5 in the gas phase.

The properties of Lewis acidity have been used in various applications, from organic synthesis to every form of catalysis to polymers. The possibilities are seemingly endless. Recent developments in the field of Lewis acidity are exciting. Mainly the upcoming fields of frustrated Lewis pairs (FLP) and electrophilic phosphonium cations (EPCs) and polymers with build in Lewis acidity, either through Lewis acidic sites or even through the whole polymer being build up from nitrogen or boron-phosphorus species are rising in popularity.

Since Lewis acids as a whole is too big of a subject to cover, some highlights out of the scientific field are chosen and discussed briefly in this thesis.

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1 - Introduction

1.1 History

Lewis acidity and basicity is an old concept, first introduced via a monograph in the American Chemical Society in 1923 by G.N. Lewis entitled “Valence and the Structure of Atoms and Molecules”.1,2_{In this monograph Lewis proposed a definition for acids and bases: “An acid is a}

substance which gives off the cation or combines with the anion of the solvent; a base is a substance which gives off the anion or combines with the cation of the solvent.” This was the first definition of a Lewis acid/base. Although the acid-base properties of a species are obviously influenced by a solvent, logically the molecular structure should be the main cause of its acidity or basicity. The work of Lewis had initially little impact on inorganic chemistry, there was virtually no mention of Lewis acidity in the first 15 years after publishing the monograph. Partially due to the imaginary boundaries between organic- and inorganic chemistry, partially because Lewis did nothing to promote his idea besides mentioning it in the monograph, but also the nature of the chemical bond itself was debated heavily during this period.2

Lewis’s view of the chemical bond and his acid-base definitions became better understood in the 1930’s mainly through the work of Pauling and Mulliken.3_{Molecular Orbital theory (MO) and Valence}

Bond theory (VB) supported the original ideas of Lewis in a simple AB system with electron wave functions.

Lewis returned to the topic of acids and bases in 1938.4_{He stated the four criteria for acid-base}

systems:

1) When an acid and a base react, the process of neutralization is rapid 2) An acid or a base will replace a weaker acid or base from its compounds 3) Acids and bases may be titrated against one another by use of indicators 4) Both acids and bases are able to act as a catalyst

After these remarks he showed that experimental acidic behavior was not confined to only the proton alone but was in general exhibited by electron-pair acceptors. He reasoned that these acids where first overlooked due to the rapid neutralization reaction that occurs in the most common solvent, water.

1.2 Definition of a Lewis Acid/Base

The definition of a Lewis acid is in its simplest sense an electron pair acceptor and a Lewis base is an electron pair donor. In this sense all organic reactions except radical or pericyclic reactions can be described as Lewis acid-base reaction. An electrophile is in this case electron seeking and thus the Lewis acid part, while the nucleophile wants to donate its electron pair, making it the Lewis base. In this context Lewis acid and electrophile are synonyms and Lewis base and nucleophile can be seen as synonyms as well. Although nucleophile and electrophile are more commonly used nomenclature while discussing reactivity and kinetics.

(1.1) The classic acid base neutralization reaction is shown in equation 1.1, this neutralization reaction is the most basic principle of an acid base reaction. The reason that this neutralization reaction works can be explained using various theories.

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1.2.1 Brønsted and Lewis acids

Brønsted-Lowry acid-base theory and Lewis acid-base theory are distinct theories but complementary. The definition of the Brønsted-Lowry acid-base theory is based upon removal or transfer of a hydrogen ion (H+_{) from the acid and its addition to the base. So, a Brønsted-Lowry acid is}

a chemical species being able to lose or donate a H+_{but it needs another substance (Brønsted-Lowry}

base) being able to gain or accept the transferred H+_{. Some Brønsted-Lowry acid examples include}

HCl, CH3-COOH and NH4+ with their respective base H2O visualized in equations (1.2-4).

+

¿

−

¿

+

H

₃

O

¿

HCl+H

2

O →Cl

¿ (1.2)

+

¿

−

¿

+

H

₃

O

¿

C H

₃

COOH +H

₂

O→ C H

₃

CO O

¿ (1.3)

+

¿

+

¿

+

H

₂

O→ N H

₃

+

H

₃

O

¿

N H

₄¿ (1.4)

As aforementioned Lewis acids are electron pair acceptors and is thus a molecule with a localized empty atomic orbital or a molecular orbital of low energy. This LUMO of the molecule can accommodate a pair of electrons and can in the process form an adduct with the electron donating species (base).

The restriction of a H+_{donor for Brønsted-Lowry acid is the key difference between both theories}

making trigonal planar species with an empty orbital, such as BF3 and AlCl3 or even metal cations,

such as Li+_{and Mg}2+_{Lewis acidic but not Bronsted-Lowry acidic.}

1.3 Physical Organic Chemistry principles

Before we dive deep into the countless possibilities, applications and a measure for Lewis acidity first we’ll discuss the basics quick. Every molecule is made up out of nuclei and electrons and bonds are made by overlap of their atomic orbitals. The Bohr Model predicts the shapes of these atomic orbitals using the Schrödinger equation which are obtained from quantum mechanics (Figure 1).3

Figure 1 Orbital shapes of the s and p orbital of carbon, with a more realistic representation by quantum mechanical calculations on the right. 3

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Pauling and Mulliken both calculated the electronegativities of atoms and the results are shown in table I below.

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Table I Electronegativities according to the scales of Pauling and Mulliken.3

Atom Pauling Mulliken

H 2.1 3.01 B 2.0 1.83 C 2.5 2.67 N 3.0 3.08 O 3.5 3.22 F 4.0 4.44 Cl 3.0 3.54 Br 2.8 3.24 I 2.5 2.88 Li 1.0 1.28 Na 0.9 1.21 K 0.8 1.03 Mg 1.2 1.63 Ca 1.0 1.30 Al 1.5 1.37 Si 1.8 2.03 P 2.1 2.39 S 2.5 2.65

The electronegativity of atoms could lead to bond polarization, in the case of a C-O bond this would lead to δ+ C and a δ- O, since oxygen has a higher electron affinity. The localization of a charge can be inductive as shown in the C-O bond, but delocalization of the electronic charge over the molecule can be visualized using Lewis resonance structures.

1.3.1 Electronic Structure

Molecular Orbital theory could help in understanding the patterns that are observed. To start with the basics, three forces are considered in the interaction energy (Ei) of a system shown in equation 1.5.

E

_i

=

E

_core

+

E

_ES

+

E

_overlap

(1.5)

As aforementioned the electrophilic part is correlated with Lewis acidity and the nucleophilic part with Lewis basicity. Ecore is the energy of the electrostatic repulsion between the electron clouds of

the two molecules, which is a positive destabilizing term. Both EES and Eoverlap are attractive and

stabilizing terms, EES is the term for electrostatic attraction between a positive and a negative charge.

Lastly, the term Eoverlap is related to the overlap of the electrophilic and nucleophilic orbitals. The

Eoverlap term explains that the interaction between the molecules is directly proportional to the size of

the electrophilic and nucleophilic orbitals but is inversely proportional to the initial energy separation between them. This correlation explains the reactivity interaction of soft nucleophiles with soft electrophiles. The preferential hard acid and base interaction is mostly controlled by the electrostatic term. This term is relatively large when charges are highly localized on the acids and bases.

Figure 2 Acceptable resonance structures.3

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1.3.2 Structure & Sterics

A molecule always seeks the lowest possible energy state, so in simple terms this means that the stability of a molecule is dependent on its Gibbs free energy according to the well-known formula (1.6).

∆ G °=∆ H °−T ∆ S °

(1.6)

Comparing molecules explains what state is preferred over another state. The difference in Gibbs free energy between two different chemical states determines the equilibrium. This is interesting of course but what makes a structure stable and another structure unstable?

Stability of a structure is not clearly defined, because it could be stable for a minute or a year or for an eternity, but a stable molecule is a molecule which always has a lower internal energy compared to a reference system. Likewise, what makes a structure reactive or unstable? The same answer can be applied, where a molecule is less stable relative to a reference molecule. This is the thermodynamic approach of stability of a molecule, but Ingold suggested another term, persistent. A persistent molecule simply means a long-lived molecule more closely tied to the kinetics with high activation barriers.

Thus far only the innate relative stability of structures was mentioned, but a key concept in the stability of molecules are the electronic and steric effects. Conjugation of π-molecular orbitals result in a delocalization of the sp2_{hybridized bond over a molecule in turn stabilizing the π-system. A fully}

conjugated hydrocarbon or heterocycle π-system complying with the 4n+2 Hückel’s rule is considered an aromatic system. An aromatic system is more stable than expected when compared to similar structures. Since acidity is about the stability of charge, structural features like sterics and electronics that can influence charge are thus crucial in the understanding and predictability of acidity.

1.3.3 Orbital effects

Electronegative elements like F, O and N generally lower the energies of all MOs. In addition, these hetero atoms introduce lone pair MOs, filled orbitals with very little bonding character. These lone pair orbitals also tend to be localized on the hetero atom presenting a large electron density on that atom.

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Figure 3: Simple representation of a molecular orbital diagram where a donor orbital and an accepting orbital overlap and lowering the overall bonding energy. 3

This interaction can be visualized using a simple molecular orbital diagram, in this diagram visualized above a low-lying donor orbital is shown mixing with an acceptor orbital. The energy difference between the donor orbital and the binding orbital is low, which gives the bond more character of the donor.

1.3.4 Reactivity

The reactivity of molecules is per definition relative. Predicting reactivity is one of the strongest tools a chemist could use, deciphering a reaction mechanism to explain the outcome of that reaction could be seen as a vital skill. Although reaction mechanisms can be well understood and accepted one should always see them as models and not as absolute truth. The reactivity of molecules can be split up in two fundamental parts, namely thermodynamics and kinetics.

1.3.4.1 Thermodynamics

Thermodynamics is a hugely important aspect of reactivity, activity and chemistry in general. Thermodynamics is based on the energy state and the change in free energy between different states of molecules measured in Gibbs free energy (ΔG). When the ΔG is negative the reaction is exothermic. This ΔG only tells us if a spontaneous change would occur but the time domain of this occurrence is determined by the kinetics of the reaction.

ΔG=ΔH −TΔS (1.7)

As shown in equation 1.7, ΔG is dependent on an enthalpy (ΔH) and entropy (ΔS) factor. ΔH is measured in kcal mol-1_{and ΔS is measured in cal mol}-1_K-1_{. Therefore the entropy factor is multiplied}

by the absolute temperature.

Ultimately thermodynamics are based on the intrinsic stability of a reactant, intermediate or product what ultimately determines its reactivity. The stability of cations and anions is thus strongly correlated with thermodynamics and acids and bases in general. Although the stability of ionic structures is hard to predict due to effects of polarity, nucleophilicity, hydrogen bonding, solvent, possible counter ions etc.

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1.3.4.2 Kinetics

Kinetics is the study which is concerned with the timescale into which one molecule transforms into another. Controlling a reaction and measuring how fast products form as a function of concentrations, temperature and various more variables provides the experimental kinetic rates of a reaction. For a molecule to transform into another molecule it needs to overcome an energy barrier. This energy barrier can be visualized by an energy surface with multiple valleys and mountains with crucial saddle points. These saddle points are the lowest energy transition barriers for the molecules, the lower saddle points are kinetically favored although these might not lead to the lowest possible energy state of the system. This is the crucial difference between thermodynamics and kinetics.

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2 - Lewis Acid/Base Strength

The strength of a Brønsted acid or base is determined by the stability of its conjugate base or acid. The stability of the conjugate species can be influenced by a multitude of physical organic principles also covered in this chapter. From electronic structure to conjugation and from sterics to orbital effects within the molecule or even solvation effects that help stabilize the buildup of charge. Taking this into account it is difficult to pinpoint what effect has the most influence on the acidity/basicity of a molecule. But the bottom line is that the thermodynamically most stable conjugate form is the strongest acid or base.

To create an extreme acidic environment, all the above-mentioned principles play a role. A so-called super acid is a mixture of a strong Lewis acid with a Brønsted acid creating some of the strongest proton donating solutions possible. A great example is one of the strongest Brønsted super acids namely fluoroantimonic acid. The combination of the strong Brønsted acid HF with the strong Lewis acid SbF5 in a stoichiometric ratio of 2:1 resulting in the super acid. This creates a medium of

non-nucleophilic and non-coordinating anions from the Lewis acids. These solutions are then possible to protonate benzene or even alkanes, which means that benzene or the alkane is in fact a stronger base than the made anions.5

+

¿

−

¿

+

H

₂

F

¿

Sb F

₅

+2 HF → Sb F

₆¿ (2.1)

2.1 Methods used to measure Lewis acidity

Determining the acidity of a certain compound can be of immense value due to the ability of making well thought out decisions and not wasting valuable time and money. The fluoride ion affinity (FIA) was first mentioned by Haartz and McDaniel in 1973 and is a method for determining relative Lewis acidities of certain compounds. This method is based upon the principal of enthalpies of formation and the fluoride affinities of these compounds.6

To determine the Lewis acidity of molecules, Stephan and co-workers have put forward the Global Electrophilicity Index (GEI) as a metric for Lewis acidity which was first introduced in 1999. 7_{The GEI}

is general, quantitative and base independent and has been benchmarked against the established FIA.

Between the FIA and the GEI some more techniques have been used to determine the Lewis acidity of compounds. Nearly all of them find basis in the electronegativity of the elements as basic principle. Pearson described in his paper in 1986 the correlation between absolute electronegativity and hardness with molecular orbital theory.8_{The hardness of a molecule is defined as twice the}

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energy gap between the highest occupied molecular orbital and the lowest unoccupied molecular orbital. Hard molecules thus have a large HOMO-LUMO gap while soft molecules have a small HOMO-LUMO gap. He proposed with this method correlations between chemical behavior, visible-UV absorption, optical polarizability, ionization potentials and electron affinities. For example, in quantum theory optical polarizability results from a mixing of suitable excited state wavefunctions with the ground state wavefunction. The mixing coefficient is inversely proportional to the excitation energy from the ground to the excited state. This means that a small HOMO-LUMO gap automatically means small excitation energies to the manifold of the excited states. Therefore, softer molecules with a smaller gap will be more polarizable than hard molecules. High polarizability is one of the most characteristic property of soft acids and bases.

A few years later in 1989 Brown and Skowron further discussed the theory of quantifying Lewis acid strength.9_{Their approach was to explore two different methods, namely Allen’s free atom}

spectroscopic electronegativity as a measure of Lewis acidity and Brown’s scale of Lewis acid strength derived from the coordination numbers observed in solids.

Allen’s scale of free-atom spectroscopic electronegativity (χ) of main group elements is defined as the average energy of the valence shell s- and p-electrons shown in equation 2.2 below.

χ=

n

s

e

s

+

n

p

e

p

n

s

+

n

p

(2.2)

In the formula ns and np are the numbers and es and ep are the spectroscopic energies for the s and p

electrons, respectively.

Brown’s scale of Lewis acid strength (Sa) is based on the principle of the oxidation state of the cation (V) and the average coordination numbers (Nt) to oxygen observed in a large sample of compounds

defined by the formula 2.3.

Sa=

V

N

_t

(2.3)

As defined in equation 2.3 the Lewis acid strength in a solid is thus the Pauling bond strength averaged over all the compound in which the cation appears. Brown and Skowron showed a clear correlation between both methods through a figure where both methods were tried out.

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Figure 5 Spectroscopically measured electronegativity correlated with Lewis acid strength in valence units for main group elements. circles represent cations in their highest oxidation state: (+) cations with stereo active lone pairs and (x)

cations with non-stereo active lone pairs. 9

Although the correlation between both methods is interesting to see, the actual usefulness of both methods is questionable due to the limited reach of only solid Lewis acidity and better-defined methods.

Some years later around 1996 Beckett et al. proposed a convenient NMR method for Lewis acidity at boron centers, also now known as the Gutmann-Beckett method.10_{The change in}31_{P NMR chemical}

shift of triethylphosphine oxide (TPO) on coordination to the Lewis acid in question is measured. The reasons that the 31_{P NMR is measured instead of the}11_{B NMR of the Lewis acid itself is simply}

because the upfield 11_{B NMR shifts upon coordination with a Lewis base are variable and often}

broadened due to chemical exchange, giving limited information about the Lewis acidity of the boron center. The more obvious reason is that Lewis acids that don’t contain a boron center can still be measured. In contrast, Gutmann and co-workers derived a quantitative parameter (acceptor number, AN) based on the 31_{P NMR of TPO. Electrophilic interaction which lead to de-shielding of the}

phosphorus atoms by inductive effects involving electron donation of oxygen to the solution. Gutmann arbitrarily chose a solvent scale with fixed points in hexane (AN=0) and SbCl5 (AN=100). In

general terms, the AN number depends upon how well the adjacent atoms bound to the Lewis acid center competes against the oxygen donor atom of the TPO for the acceptor orbital.11

A conceptually similar method to the Gutman-Beckett method is the Childs method. The Childs method is based on crotonaldehyde as the base, and the change in 1_{H NMR chemical shift of 3H is}

probed. Recently, Hilt and Nödling used 2_{H NMR spectroscopy to assess the change in}

para-deuterium signal of d5-pyridine with a change in Lewis acid (Figure 6).12

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While both the Gutmann-Beckett and the Childs method and some variations work fine for determining the strength of a Lewis acid and are definitely useful, neither give predictive Lewis acid strength. This ability to predict Lewis acidity of for instance newly synthesized compounds could be of immense value. Both as a guideline for synthesis to know beforehand which Lewis acid is preferable over another in synthesis and expectation value that could guide design of new systems. The briefly before mentioned methods fluoride ion affinity (FIA) and global electrophilicity index (GEI) do provide a predictive feature for the Lewis acidity strength. Christe et al. published a paper in 2000 introducing his method (FIA).13_{FIA is now commonly used with modern electronic structure}

computations, by calculating the change in enthalpy accompanying coordination of a fluoride ion to the Lewis acid in question. Fluoride is chosen in this technique for its high basicity, small size and its willingness to form adducts with a wide variety of Lewis acids.

Since only some fluoride ion affinities had been estimated experimentally and with large discrepancies due to the use of various techniques, Christe and coworkers developed an important self-consistent basic set of fluoride ion affinities.13_{Using the theoretical fluoride affinity calculations}

at the MP2/PDZ level of theory for the reaction depicted in equation 2.4.

−

¿

−

¿

+

A → CO F

₂

+

A F

¿

C F

₃

O

¿

(2.4)

The relative fluoride ion affinity scale is based on this this reaction with an experimentally known value of 49.9 kcal/mol. The relative fluoride ion affinities in kcal/mol are divided by 10 to obtain a convenient scale resulting in a set of numbers as a quantitative pF-_{scale for Lewis acidity.}

Table II Abbreviated pF-_{scale for various Lewis acidic compounds.}13

Compound pF -SbF5 12.03 AlF3 11.50 AlCl3 11.46 PF5 9.49 BF3 8.31 SiF4 7.35 SF4 5.67 PF3 4.49 HF 3.68

The pF-_{for multiple Lewis acidic compounds are measured and shown in Table II and shows the}

expected trend for acid strength experimentally observed, SbF5 > AlF3 > BF3 > PF3 > HF. The higher the

observed value the greater the relative strength. There are some side notes that need to be mentioned about this method. The given values are all calculated for the free gaseous molecules, the values of pF-_{for polymeric molecules like AlF}

3 are actually smaller than the given values in the table

since no correction for their association energies is involved. The stability of a fluoride complex can be predicted from the pF-_{value and should be above ~3.5.}

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The FIA is an interesting tool for chemists, but has inherently some flaws that cannot be overlooked or overcome easily, mostly because it is based on the very hard Lewis base F-_{and mostly focused on}

the stability of complexes and more accurately describe the fluoridophillicity of the molecules. Some other techniques have been put forward using a hydride or chloride to correlate with the Lewis acid strength, but both have the same inherent problem that is observed by the FIA as well.

The aforementioned GEI was first introduced in 1999 by Parr et al.14_{and suggested by Stephan and}

co-workers as a useful metric for Lewis acidity. GEI, usually abbreviated with ω, is a measure of the ability of a molecule to take up electrons, defined in the formula 2.5.

ω=µ

2

/2 η= χ

2

/

2η

(2.5)

µ is defined as chemical potential while η is the chemical hardness and χ is the electronegativity. According to the definition of Mulliken µ is the negative of χ resulting in both formulas for ω.15_This

formula agrees with the basic chemical principles that more electronegative and softer elements will have a higher ratio of taking up electrons. A good thing to mention is that the GEI is used to assess the electrophilicity of molecules, where electrophilicity is a kinetic phenomenon. This is related but different from Lewis acidity which is a thermodynamic concept.

The main practical advantage of the GEI over the FIA for calculating relative strengths of Lewis acids is the theoretical calculations. Where the FIA uses the reaction shown in (2.4) as a baseline to calculate the relative stabilities of the other Lewis acid-base adducts resulting in more and longer calculations/computing power. The values used for calculating the GEI are obtained solely from the HOMO and LUMO of the optimized Lewis acid in question using some simple formulas 2.6/7.

µ=

1 ₂

(

E

HOMO

+

E

LUMO

)

(2.6)

η=

(

E

LUMO

−

E

HOMO

)

(2.7)

Stephan correlated some calculated values for boranes with both the established FIA and the GEI method. The boranes used and the results that were observed are both seen in Figure 7.

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Figure 7 Boranes explored in this study coded 1-4 with corresponding GEI and FIA values.7

The two methods are compared using different units (ω is given in eV and FIA in kJ/mol by convention), and a direct comparison of absolute values is meaningless as different properties are being measured. But the qualitative trends between both methods, given by the steepness of the curve, holds true for both GEI and FIA. This trend holds mostly true for trityl cations and phosphonium cations, which are also calculated in the paper. A more interesting correlation is that of the sulfoxonium species explored in this study with the results given in the figure below.

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Figure 8 Sulfoxonium species explored in this study coded 5-9 with corresponding GEI and FIA values.7

In Figure 8 a clear discrepancy between compound 9 is shown. Using the FIA calculations this molecule is supposedly equally acidic as compound 5 while GEI calculation predict compound 9 to be some order of magnitude more acidic than compound 5. Using the Gutmann-Beckett and the Childs methods mentioned before to correlate the calculated GEI and FIA values some correlation can observed. The FIA strongly correlate with the Gutmann-Beckett method for measuring the relative Lewis acidity of compounds, while GEI shows more correlation with the Childs method. Since FIA uses a fluoride to determine Lewis acidity, which is a hard Lewis base it is not surprising that it correlates better with the somewhat harder base used in the Gutmann-Beckett method. Those results however indicate that GEI could be used as a more accurate measuring tool for softer Lewis acids. Although no clear explanation is given for the observed correlation.

Recommendations from the author would be to use FIA for harder Lewis acids and the GEI for softer Lewis acids, but both techniques mostly focus on the electronic part of Lewis acidity with little to no consideration about physical chemical properties like sterics or molecular orbitals. This is understandable since both techniques are mainly used for a quick scan to determine Lewis acidity and no absolute clear-cut absolute ranking of molecules similar to the pH scale that is used for Brønsted acids and bases. However, since GEI can be performed using data gathered by calculating the FIA it can still be seen as a quick scan as an indicator of Lewis acid strength.

Baumgartner, Caputo and co-workers recently suggested a new correlation of chromaticity data with Lewis acid strength.16_{The authors mention a fluorescence based method based on the stable}

dithieno[3,2-b:2’,3’-d]phosphole. The interactions dithienophosphole has with Lewis acids is displayed in Figure 10. This highly luminescent and easily tunable fluorophore was first introduced in

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2004.17_{Fluorescent probes have since found successful applications for the sensing of small}

molecules, biological species, anions (F-_{and CN}-_{) and in materials science.}

Figure 9 Structure of the dithienophosphole after Lewis acid coordination on the phosphoryl group.16

The main benefits of using fluorescence in the signal-transduction process are derived from high sensitivity, allowing for low detection limits around the ppm range, which can exhibit a high degree of analyte specificity and providing the distinct ability of often visible and easily detectable change in the probe’s luminescence properties. Additionally, experiments can easily be done under inert atmosphere to account for the highly sensitive Lewis acids.

Figure 10 The luminescence properties of dithienophosphole shown in vials.17

The authors showed that the optical properties can be effectively fine-tuned in various ways.17

However, extensions of the conjugated scaffold leads to a range of emission colors that span the full optical spectrum, the substitution pattern of the phosphorus also has a noticeable impact on the optical properties of the system. The latter is the result of the phosphole-typical σ*-π* interaction between the exocyclic substituents and the conjugated system which alters the LUMO energy level as a function of the polarity of the exocyclic bond(s). This means that in general, the more electronegative the exocyclic bonding partner is, the larger the contribution of E to the σ-orbital that will concurrently lead to a larger contribution of the phosphorus atom to the σ*-orbital, which creates more overlap with the π*-system (Figure 11). This ultimately translates to a lowered LUMO, which ultimately translates to a red shifted emission that can be tuned by the electronegativity (EN) of the exocyclic substituent.

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Figure 11 Schematic overview of the interactions dithienophospholes with Lewis acids, where in (A) the impact of Lewis acid strength on P-O bond is shown and in (B) the impact of the Lewis acid strength on the energy levels of σ*-orbital and

its interaction with the π*-system leading to overall altering of LUMO energy is shown.16

This simple and powerful naked-eye litmus test for Lewis acid strength might be able to resolve some previous hiccups before mentioned. The primary goal was a proof of concept study wherein they reported the detailed fluorescence and chromaticity responses to the dithienophosphole system in the presence of a series of important Lewis acids to showcase the available Lewis-acidity range of their method. To set a foundation for the fluorescence Lewis acidity (FLA) test, three dithienophosphole oxide probes with different emission wavelengths were used shown in Figure 12.

Figure 12 Dithienophosphole oxide probes utilized is creating FLA scales.16

B(C6F5)3 is selected as a representative Lewis acid to test. The reaction of 10 with one equivalent of

B(C6F5)3 in CH2Cl2 at room temperature resulted in an immediate shift in the emission color of the

reaction mixture from sky blue to aquamarine (under UV-Lamp). Furthermore, multinuclear NMR analysis further indicates adduct formation between the phosphoryl group and the borane. Most notably, the 31_P{1_{H} NMR resonance shifted downfield to 30.2 from 18.2 ppm. The}11_{B NMR spectrum}

showed a resonance at -1.31 ppm, and the 19_{F NMR spectrum showed three resonances for the}

ortho-, para-, and meta-fluorine atoms of the pentafluorophenyl rings at -133.2, -158.0 and -164.3 ppm, respectively, further supporting adduct formation.

The coordination of the borane species to 10 was definitively proven via single-crystal X-ray crystallography (Figure 13). Interesting observations from the crystal structure include the P-O-B angle of 130.59(13)ο_{and the π-stacking observed between one of the pentafluorophenyl rings and}

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Figure 13 Solid state structure of the adduct 10-B(C6F5)3 (50% probability level).16

To illustrate the value of the FLA a broad scope of Lewis acids ranging from transition metal- to main group-Lewis acids. The Lewis acids were all tested with toluene as solvent. AlCl3 and BF3 will be kept a

close eye on, since both compounds were also mentioned in the previously FIA range of Lewis acidity.

Figure 14 (A) Normalized emission spectra for the FLA with 10 (B) Magnified area between 430 and 563 nm.16 In Figure 14 the normalized emission data spectra for the FLA with 10 are shown. Differences in emission between adducts of weak and strong Lewis acids can easily been seen even by the naked eye. The weak Lewis acid BPh3 generates a FLA that has a red-shifted emission of 6nm, whereas the

FLA with [Et3Si][B(C6F5)4] exhibits a bathochromic shift of 69 nm (Figure 10). The results of all adducts

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Figure 15 Chromaticity trends of adducts of all Lewis acids where A=10, B=11, C=12.16

The trends found in Figure 15 are used for each individual Lewis acid and the true Lewis acidity is calculated giving a rating that is visualized in Figure 16.

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Figure 16 Scale of the strength of all the tested Lewis acids.16

Although not perfect the new method of Lewis acidity rating with three simple fluorescence measurements could provide the next step in the development of a true Lewis acidity scale.

2.2 Physical Organic Principles of Lewis Acidity

An interesting paper published in 2011 by Ashley and co-workers described the separation of electrophilicity and Lewis acidity.18_{The widely used and well-known tris(pentafluorophenyl)borane}

(B(C6F5)3) Lewis acid was compared with a somewhat similar range of molecules with

pentachlorophenyl groups added to the borane. Tris(pentafluorophenyl)borane was first synthesized in 1963 by Massey and his colleagues,19_{immediately noticing its tendency to form strong adducts}

with phosphines, ammonia and ethers. B(C6F5)3 has found numerous applications in organic and

inorganic chemistry ranging from silylation of alcohols, hydrosilylation of ketones and imines, reductive cleavage of ethers and alcohols to synthesis of weakly coordinating anions, anion binding and activation of transition metal mediated α–olefin polymerizations.20–24_{More recently B(C}

6F5)3 is

used in the field of frustrated Lewis acid base pairs. All these attributes are related to B(C6F5)3 being a

strong Lewis acid, which has been measured to be intermediate of BF3 and BCl3.25 However, the big

advantage of B(C6F5)3 over BF3 and BCl3 is its physical properties of being a thermally robust solid and

water resistance lending itself to ease of handling. The stability, even at elevated temperatures of B(C6F5)3 stays robust, combined with some steric bulk making it an ideal boron-based Lewis acid.22,26

The comparison of the B(C6F5)3 Lewis acid with an increasing number of substituted C6Cl5 phenyl rings

to the borane center at first glance might not seem as interesting. (Figure 17) A chloride is one row below fluoride and looking at the electronegativity trend in the periodic table leading to the conclusion that χpauling of fluoride being 3.98 while χpauling of chloride is 3.16.27 Based solely on

inductive effects, one would expect Lewis acidity ranking of B(C6F5)3 > B(C6F5)2(C6Cl5) > B(C6F5)(C6Cl5)2 >

B(C6Cl5)3. However, mesomeric effects are very important here.

Figure 17 Space fill diagram of (a) 1 substituted ring of chlorides (b) 2 substituted rings with chloride and (c) 3 substituted rings with chloride. Cl atoms orange, F atoms green, B atoms pink.18

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Mesomeric donation from the ortho- and para-F lone pairs is particularly effective due to the molecular orbital 2p-2p overlap with the adjacent carbon.28_{This results in significant π-back donation}

in the aromatic π-system and subsequently into the acceptor orbital of the boron which decreases relative Lewis acidity. While chloride is not as electronegative as fluoride, the mesomeric π-back donation into the aromatic system is substantially lower due to the weaker 3p-2p overlap π-overlap. This effect resulted in a net increase in electron deficiency at the boron for pentachlorophenyl groups (C6Cl5) over C6F5 groups.

Interestingly, Ashley tested the electron deficiency at the boron center for all the B-(C6F5)3-n-(C6Cl5)n

(n=1-3) species with electrochemical and theoretical studies using DFT calculations, showing that C6Cl5 is a more electron withdrawing substituent relative to C6F5.18 However, a decrease in Lewis

acidity is measured from using C6Cl5 over C6F5 using both the Gutmann-Beckett and the Childs

method mentioned earlier in this chapter. This is an interesting finding because it is in contrast to the believe that Lewis acidity is only a measure of electron pair uptake ability. While electrochemistry provides a physiochemical measure of electron affinity at the boron center in these compounds it completely neglects the steric cost of B sp2_-sp3_{hybridization upon coordination with a Lewis base,}

which is a more important factor for bulkier boranes. However, The Gutmann-Beckett/Childs methods do incorporate both the electronic and steric factors in their measurement for Lewis acidity giving it a more reliable indication of “real Lewis acidic strength”.

These results may be confusing at first glance but can be rationalized by realizing that Lewis acidity is clearly not only a measure of electronics, but a more complex phenomenon involving multiple physical organic basic principles. The affinity of a Lewis acid to a Lewis base thus strongly depends upon the interplay of attractive forces (electrostatic, covalent, dispersive) and repulsive interactions.5

A schematic overview of all the interactions is shown in Figure 18.

Figure 18 The contributions of Lewis acid-base pair formation in schematic overview depicted.5

Different Lewis acid affinities toward different Lewis bases reveal that the electronic interactions between both Lewis acid and base are not the only contributing factor and on its own not sufficient enough to determine Lewis acidity. Molecular orbital overlap is an important feature mostly for softer Lewis acid-base adducts, while the electrostatic interaction is of greater importance in harder Lewis acid-base adducts. Combining interactions into one formula provides us with formula 2.8.

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ΔH =E

_LA

E

_LB

+

C

_LA

C

_LB

+

W

(2.8)

E are the electrostatic contributions of both the Lewis acid and the base, while C is used for the covalent contribution and W is a measure of steric interaction.

2.3 Super Lewis Acids

Lewis acidity/basicity is a complex phenomenon that arises from multiple basic physical organic principles mentioned before in this chapter. As shown before multiple effects influence the strength of the Lewis acid for example: molecular orbital overlap, steric hindrance, aromaticity, mesomeric effects and inductive effects just to name a few. The term super Lewis acids or simply Lewis superacid (LSA) for short is combining the effects of all the elementary physical organic chemistry principles to increase the acidity to the extreme.

Before we discuss ideal conditions and physical organic chemistry principles to create the most acidic Lewis acid possible, we first need to define what makes a simple Lewis acid a LSA. Should a LSA be stronger than antimony pentafluoride (SbF5) or should it provide superb catalytic activity like for

instance the zinc(II) Lewis acidic center at physiological pH values in biological systems?29,30_{Or is the}

so called “Lewis super acid” describing the peculiar behavior of SiF3+ in the gas phase described by

Speranza et al.?31,32_{While Dunach and his coworkers called metal triflimides and triflates LSA as they}

are arrived from Brønsted super acids HNTf2 or HOTf.33 Olah defined LSAs as those acids that are

stronger than anhydrous aluminum trichloride.34_{Nonetheless, in all the above-mentioned cases,}

neither the respective Lewis base nor the standard state was specified. Krossing gave the most coherent definition: “Molecular Lewis acids, which are stronger than monomeric SbF5 in the gas

phase are Lewis Superacids.” 30

Before mentioned FIA is mostly used to compare different Lewis acidities with one another. Next to some drawbacks already mentioned before, it has another mayor drawback which is its one-dimensional value which is particularly impactful for softer Lewis bases like methide or hydride, where differing orders exist. For example, B(C6F5)3 is not a LSA judged by the FIA, whereas Al(C6F5)3 is.

Nonetheless, B(C6F5)3 has a higher hydride ion affinity (HIA) than Al(C6F5)3 making it more susceptible

for the relatively softer hydride Lewis base.35_{A fairer representation of LSA might be a combination}

of both the FIA and the HIA to give a complete picture. Greb proposed a more fitting definition for LSA where he consequently put more appreciation in the special role of B(C6F5)3 in the recent years.

Making the definition as follows:

“Molecular Lewis acids that have a larger FIA than SbF5 in the gas phase are Lewis super acids (LSA).

Molecular Lewis acids that have a larger HIA than B(C6F5)3 in the gas phase are soft Lewis super acids

(sLSA).” 5

Using this definition for sLSA permits a better consideration for softer Lewis acidity features. sLSA might stimulate research in the molecular orbital overlap direction and orbital-controlled bond activations, which seems to be the predominant factor in sLSA. 36–38

Already shortly before mentioned in the previous chapter are the super acids where a Lewis acid and a Brønsted acid combine to form an extremely potent acid which is even able to protonate benzene. These species are called Brønsted super acids. Most strong Brønsted acids are based on the principle

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of weakly coordinating anions (WCA). The stronger the Lewis acid the more stable the WCA proves towards decomposition.

So far, the possibilities sound interesting and fascinating, but for application purposes a LSA has to have some features to be useful:

1.

A high thermal stability of the Lewis acid with kinetically or thermodynamically blocked decomposition pathways.

2.

Ligands are strongly bound and inert making side reactions by ligand scrambling or solvolysis thermodynamically or kinetically prohibited.

3.

The Lewis acid is water-tolerant and decomposition upon liberation with HF does not occur.

4.

The Lewis acid is non-reducing (high ionization potential and large HOMO/LUMO gap) and

non-oxidizing (low electron affinity).

5.

The Lewis acid is cheap and should be easily synthesized from commercially available precursors.

6.

The Lewis acid is spatially well defined, and it can be easily adjusted making its steric properties variable.

7.

The Lewis acid is protected against poly-/oligomerization processes, as these often result in poorly defined active species with diminished Lewis acidity.

8.

Solubility as well as solvent tolerance is high for non-polar as well as for polar solvents. Although most LSAs won’t comply with all the rules set above, the more above-mentioned features it has it gains essential advantages compared to usual metal halides. In turn making the LSA better and broader for potential use.5

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3 – Lewis Acid applications

Lewis acids are interesting molecules with interesting properties. As already before-mentioned it is capable of accepting negative charge due to a variety of physical organic effects. In this section a variety of applications for Lewis acids will be discussed. Since the field and use of Lewis acids as a whole is too broad, not every single application where Lewis acids are involved will be mentioned. The applications of Lewis acids that will be discussed are catalysis, frustrated Lewis acid base pairs (FLPs), electrophilic phosphonium cations (EPCs), polymers and Lewis super acids (LSA).

Catalysis and polymers as Lewis acid applications are chosen because of the huge influence of both fields on our current way of life. The other discussed applications (FLPs, EPCs, LSA) are chosen because of their recent developments being relatively new fields where Lewis acidity is playing a vital role.

3.1 Catalysis

While talking about the applications of Lewis acids the first thing that will come to mind is of course catalysis. Since catalysis is an integral part of our lives and the wide field of applications of Lewis acids one would think both fields show some overlap. Centi and Perathoner already mentioned the use of acidified clays for hydrocracking in their review of catalysis in 2007, now already over 90 years ago at the start of the petrochemical industry.39_{This just to mention the long history of Lewis acids}

and catalysis as a whole.

Acid catalyzed reactions are by far the most numerous and best-studied reactions in catalysis.40_As

already mentioned it is beyond the scope of this literature thesis to mention all the uses of Lewis acids in organic synthesis or catalysis. Instead, showing more the gradual transition from environmentally non tolerable homogeneous Lewis acids toward solid and green Lewis acids for catalysis. Furthermore, some key industrial processes which depend on Lewis acidity will be mentioned as well.

The Lewis acids alkylaluminum chlorides (RnAlClm) and aluminum trichloride (AlCl3) are used for

catalytic ene cyclization.41,42_{Some examples are of the cyclization of the terminal}

methylpropenyl-substituted α,β–ynals and for the asymmetric ene reaction of chiral α–cyanovinylic sulfoxide shown in Figure 19.

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Aluminum trichloride is also used as catalyst in the rearrangement of 1-indanone oxime. This reaction was performed in -40 ο_{C in dichloromethane and yielded 91% of the corresponding aromatic}

lactam shown in Figure 20 below.43_{The use of more conventional Brønsted acids such as poly}

phosphoric and sulfuric acid could be used as catalyst for this reaction, but the yield is much lower at around 20%.

Figure 20 Rearrangement reaction catalyzed by Lewis acid.43

Alkylaluminum chlorides are typically added as a co-catalyst to the Schrock metathesis.44_{The Schrock}

reaction is a powerful synthetic reaction which can make large macrocycles by cyclization of open precursors bearing many different reactive functionalities. The Lewis acid is in this case used to enhance the activity of the ruthenium organometallic species.

Figure 21 Schrock metathesis reaction with a Lewis acid as co-catalyst.44

Another great example of a Lewis acid as a co-catalyst to enhance reactivity is in alkene polymerization reactions. The radical polymerization of acrylamide, methyl methacrylate, styrene and vinyl acetate grafted into 1,2- and 1,4-polybutadiene in benzene at 60ο_{C are all enhanced with}

the use of a small amount of AlCl3 or ZnCl2.45,46

A good example of the influence of Lewis acid strength in polymerization reactions is that of titanium acid strength in the living cationic polymerization of styrene. The polymerization of styrene is achieved in dichloromethane at temperatures between -40 and -78ο_{C by promoting the styrene-HCl}

adduct with TiCl3(OiPr) as Lewis acid. A stronger Lewis acid (TiCl4) produced low quality polystyrene

while a weaker titanium Lewis acid (TiCl2(OiPr2)) did not induce the polymerization reaction at all.47

The acetalization of C=O by alcohols or glycols and hydrolysis of the made acetals are typical acid catalyzed synthetic reactions. BF3•Et2O Is a Lewis acid which promotes acetalization of carbonylic

compounds, formation paraformaldehydes and copolymers with oxymethylene units and trans-acetalization reactions. The dimethyl acetal of (1S)-(+)-N,N-diisopropyl-10-camphorsulfonamide reacting with BF3•Et2O as catalyst and α–hydroxyacids producing a chiral 1,3-dioxolate (Figure 22).

The enolates derived from these compounds undergo reactions with high diastereoselectivity with alkyl halides. Subsequent hydrolysis results in mono- and di-substituted α-hydroxyacids with high enantiomeric excesses.48

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Figure 22 Acetal formation in organic chemistry with the aid of Lewis acids.48

The well-known Friedel-Crafts reactions in organic synthesis are another example of catalytic Lewis acidity using mainly AlCl3.49–51 The conventional reaction mechanism for the Friedel-Crafts reaction

suggests that the role of the Lewis acid is to generate the electrophilic acyl or alkyl species without interaction with the aromatic compound. While ab initio calculations suggest that one benzene carbon becomes highly nucleophilic, facilitating the attack of incipient, not fully developed electrophiles in the presence of a Lewis acid. An example of an AlCl3-catalyzed Friedel-Crafts

alkylation is that of propylene oxide reacting with benzene and alkyl derivatives at 0ο_C.52

Figure 23 Lewis acid catalyzed Friedel-Crafts reaction.52

The Mukaiyama aldol condensation reaction consist of the reaction of silyl enol ethers with ketones or aldehydes in the presence of Lewis acids creating a valuable C-C bond.51_{The strong TiCl}

4 Lewis acid

is preferred for these reactions. An example of the aldol reaction is the trimethylsilyl enol ether of acetophenone reacts with acetone in the presence of TiCl4 to give

1-phenyl-3-methyl-3-hydroxy-1-butanone with a 74% yield shown below.53

Figure 24 Aldol reaction catalyzed by Lewis acid.53

Furthermore, heterocyclic rings as for instance: small cyclic ethers, epoxides, oxetanes and ozonides having an alkoxy group on the side chain, can in the presence of a Lewis acid rearrange to give the ring expanded cyclic ethers or addition products. An example of stereoselective Lewis acid catalyzed rearrangement can be seen in aryloxyepoxy alcohols converted to the functionalized chroman diols (Figure 25).54–57

Figure 25 Stereo selective rearrangement catalyzed by Lewis acids.57

Even one of the most powerful synthetic organic chemistry C-C bond formation tools in the Diels-Alder reaction can be catalyzed via a Lewis acidic catalyst.58–61_{The Diels-Alder reaction is a pericyclic}

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6-membered ring structures. The reaction requires HOMO-LUMO overlap and inverted electronic demand for the dienophile and diene. This electronic density mismatch can be overcome with the use of Lewis acid as catalyst for the reaction. Concerning the mechanism of this catalyzed [4+2] Diels-Alder cycloaddition using Lewis acids, it has been established that dienes, being electron rich, undergo an electron abstraction by the Lewis acid to form the corresponding diene radical cation as visualized in Figure 26. This electrophilic species can be seen as the actual reaction intermediate through a chain mechanism.

Figure 26 Proposed mechanism for the Lewis acid catalyzed Diels-Alder pericyclic reaction.58–60

The mechanism proposed in Figure 26 above is further promoted by the photochemically and the radical cation promoted Diels-Alder reactions that show a similar mechanism. In its simplest form, the mechanism of the radical-cation mediated cycloaddition starts with one of the addends undergoing electron transfer to a one electron oxidant.60,62_{This redox umpolung process in turn}

activates the molecule and a cycloaddition ensues. The product is then formed by electron transfer to the cycloadduct radical cation from either another molecule of starting material or the reduced form of the initial oxidant. Saettel et al. tested the proposed mechanism with quantum chemical calculations at B3LYP/6-31G* theory level, concluding that the basic mechanism as shown in Figure 26 is on solid footing.58

More recently Nakayama et al. demonstrated the radical cation Diels alder reaction.63_{They described}

the TiO2 photocatalyzed Diels-Alder reaction between electronically mismatched reactants anethole

and 2,3-dimethyl-1,3-butadiene. (Figure 27)

Figure 27 TiO2 photocatalyzed Diels-Alder reaction.63

The authors suggested a similar mechanism as already before mentioned and compared it to the photochemically possible [2+2] addition, which has a similar radical cation mechanism as the

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proposed [4+2] Diels-Alder mechanism. (Figure 28) With optimized conditions the reaction depicted in Figure 27 proceeded with excellent yield (>90%) without any difficulty, adding steric bulk to the diene was also found to be a successful substrate (16). Product 17 was obtained from (E)-1,3-pentadiene, whereas the use of (Z)-1,3-pentadiene was unsuccessful. Trans-cycloadducts where selectively formed from trans-anethole (13), using anethole was found to give a mixture of cis-and trans-cycloadducts. This suggests that the previously mentioned stepwise pathway, with loss of stereochemistry can potentially happen by bond rotation of the anethole radical cation.

Figure 28 Schematic overview of photocatalytic cycloadditions with TiO2 as catalyst.63

Multiple reports have also shown the activity and region- and stereoselectivity of Lewis acids as catalysts for Diels-Alder cycloadditions.58,59,61_TiCl

4 and SnCl4 have been found to catalyze

stereoselectively the reaction between 2-methoxy-1,4-benzoquinones with unsymmetrically substituted stilbenes creating trans-2,3-diaryl-2,3-dihydrobenzofuran-5-ols.64

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Figure 29 Stereo selective Diels-Alder product catalyzed bySnCl4.64

Even some [2+2] cycloadditions have been reported that are catalyzed by Lewis acids. For example, methylthio-substituted allenylmethylsilane undergoes [2+2] cycloaddition with alkenes in the presence of ethylaluminum-dichloride, to afford corresponding methylenecyclobutanes,65,66

this is readily converted further to di-exo-methylenecyclobutanes by sequential oxidation of the methylthio to sulfone and 1,2-elimination.

Figure 30 [2+2] cycloaddition promoted by Lewis acid.66

Lewis acids catalyze plenty more reactions, but the before mentioned reactions give a small overview of the wide variety of reactions made possible by Lewis acids. An often more “green” and sustainable catalysis pathway is heterogeneous catalysis, where the Lewis acid is immobilized on a supportive material with high surface area like Al2O3, SiO2, graphite or zeolites to name a few, to perform the

catalysis.67_{Common Lewis acids used for heterogeneous catalysis are FeCl}

3, SbF5 and aluminum

halides. Possibilities for incorporating Lewis acids in the supporting material are physisorption, chemisorption or anchoring. The most commonly used Lewis acid for Al2O3 is AlCl3, this Lewis acid

strongly interacts with the surface of the alumina forcing the material to convert and create new active sites shown in Figure 31.

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Figure 31 Interactions of the support Al2O3 with the Lewis acid AlCl3. 67

The applications of these Lewis acid heterogeneous catalysts are as wide if not even wider than the homogeneous equivalent, from alkylation reactions to isomerization. These heterogeneous methods are preferred over the homogeneous variant, mainly due to the ease off handling, separation and “greener” reaction conditions.68

3.2 Frustrated Lewis acid base pairs

A more recent development in Lewis acidity and catalysis are frustrated Lewis acid base pairs. As before mentioned, Lewis classified acids and bases as electron-accepting and electron-donating molecules respectively. Using this simple model, one could predict that a simple donor phosphine molecule and an accepting borane molecule should react to form a P-B bonded adduct and indeed that occurs. As this reaction might be predicted and seemingly easy, once geometric constraints and/or sterically demanding ligands are added can deter the bonding between the phosphorus and borane adduct. Such a system is then a so called ‘frustrated Lewis acid-base pair’ (FLP).69_This

phenomenon was recognized early on, for example Brown et al. already showed that lutidine formed an adduct with BF3 but not with the bulkier B(CH3)3.70 Subsequently, the Wittig group in

situ-generated benzyne with the 1,2-addition reactions of PPh3/BPh3 pair and showed that the CPh3-/BPh3

pair adds to butadiene.71,72

Figure 32 Simple visualization of the frustrated Lewis acid/base adduct. 69

To give a clearer visualization Stephan has drawn a schematic overview of the frustrated Lewis acid and base pair.69_{Frustrated Lewis acid and base pairs are not per definition 2 separate molecules that}

cannot covalently bind due to steric hindrance. Several molecules with intermolecular Lewis acidity and Lewis basicity functionality have been mentioned and some examples are shown below.

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Figure 33 Some examples of intermolecular Lewis acid/base interactions.71

This acidity and basicity functionality in the same molecule are not new and is vital for life itself. As enzymes have used this functionality within the same molecule for centuries to make various reactions possible that would not be possible at the neutral pH of 7 in the body.69

Stephan et al. showed that the heterolytic splitting of hydrogen using an FLP is possible.73_This

activation of molecular hydrogen was long believed only possible with transition metal complexes. The phosphine and borane adduct is able to thermally or catalytically eliminate H2.74 The authors

proposed a mechanism for this activation which is shown in the scheme below.75

Figure 34 Proposed mechanism for hydrogen activation. 75

Figure 35 Hydrogen activation by a FLP adduct structure. 76

This mechanism was tested in 2015 by Stephan and the resulted structure is shown above.76,77

Stephan determined via the Grimme model a “side on” interaction with B and “end on” interaction with P. This orientation allows σ donation from the H-H bond to B with P donation to the H2 σ*

orbital.

But in a recently published review by Paradies the mechanism of the metal free H2 -activation was

revised.78_{Quantum chemical data suggests that the diffusion of H}

2 into the cavity of the encounter

complex leads to the weak polarization of the H-H molecule within the electric field (EF) preceding orbital interactions with the Lewis pair’s heteroatoms.76,79_{The Lewis base donates electron density}

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into the σ(H2)* orbital of the pre-polarized H2, this leads to the destabilization of the H-H bond

(electron transfer, ET process). A clear cut between the ET and the EF process is not possible to determine for FLPs, which are active in the splitting of H2. It is reasonable to assume that both

processes occur in the course of the H2-activation, although probably at different reaction

coordinates. Pre-polarization through the EF process illustrated in Figure 36 below is vital to stabilize the charge separation and to lower the σ(H2)for efficient ET of the corresponding Lewis base and

charge stabilization by the Lewis acid.

Figure 36 Proposed mechanism of H2 activation by FLPs by preparation and electric field (dashed lines). 76 Toluene solutions of stoichiometric mixtures of R3P (R=t-butyl, C6H2Me3) with B(C6F5)3 were

monitored by 31_P{1_H},1_H,11_{B and}19_{F NMR spectroscopy. No formation of the Lewis acid base pair was}

observed at a temperature range from -50 ο_{C to 25}ο_{C which is consistent with the sterically}

demanding substituents on the boron and phosphine.75_{When exposed to H}

2 at 1 atm pressure at 25 ο_{C white precipitants started to form quantitatively. NMR data of the product were consistent with}

the formulation as [R3PH][HB(C6F5)3].75

With the formation of heterolytically cleaved hydrogen-FLP adducts, possible catalytic hydrogenations can be thought of and tested.80_{The first}

report of successful hydrogenations using FLP adducts was with sterically demanding aldimines, aziridine and protected nitriles shown in Figure 37. Furthermore, it was recognized that imine substrates could act as the FLP base, and thus only needing the B(C6F5)3 as catalyst and H2 to

perform the reduction catalytically. Mechanistically these reactions proceed by the protonation of the imine and subsequent transfer of the hydride from the hydridoborate to the iminium carbon.74_{In the end}

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extended to reduce diimines, pyridyldiimines and even imine precursors for antidepressants, anticancer drugs and herbicides.

Subsequent studies showed the broadened substrate scope of FLP catalysts, where Rieger prepared intra molecular intramolecular amine-borane catalyst.81,82_{While the Erker group increased the}

substrate scope by using silyl enol ethers.83,84

3.2.1 Electrophilic phosphonium cations

The field of FLPs is emerging and gaining a lot of attention currently, this is the main reason that the applications and uses of FLPs won’t be discussed extensively, but the newly emerging Lewis acid group of electrophilic phosphonium cations (EPCs) should be mentioned (Figure 38).

Figure 38 Example of an electrophilic phosphonium cation.85

EPCs have long been known to act as Lewis acid catalysts. Their main use is in catalytic C-C bond formation. The first time EPCs were used was by Mukaiyama et al. in 1989,86_{the dicationic catalyst}

[R3POPR3]2+-[OTf]2- was able to perform Mukaiyama aldol reactions of aldehydes and silyl enol ethers

or silyl ketene acetals, promote Michael reactions of enol ethers and allylic silanes as well as Mannich-type reactions of imines silyl ketene acetals.87

Typical main group Lewis acids consists of group 13 elements, where the Lewis acidity can be modulated by appropriated substitution of the Lewis acidic center. A common way of increasing Lewis acidity is by substitution of the central atom by a strong electron withdrawing group like perfluorinated arenes. Alternatively, to electron withdrawing groups a positive charge can be incorporated on the Lewis acidic center. However, these cationic B, Al or even group 14 C and Si species show high reactivity towards nucleophiles often leading to stoichiometric formation of stable compounds. The electrophilic phosphonium salts consist of group 15 element phosphorus and in contrast to the before mentioned cations are generally less reactive to nucleophiles.

The ability of some PV_{phosphonium cations [R}

4P] + to form neutral phosphoranes R4PX is evidence for

their enhanced Lewis acidic character. In fact, such reactivity is the key step in the Wittig reaction, where the Lewis acidic phosphorus center in ylides accepts electrons from the oxygen atom of the carbonyl.88_{This is the property that is used in the development of catalysis.}

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In 2013 Stephan et al. revived this field of science with the introduction of a EPC, containing a highly electrophilic fluorophosphonium cation [(C6F5)3PF] +[B(C6F5)4]-.89 Several strongly Lewis acidic EPCs

have since been developed with a variety of applications. The focus will however be on the more recent developments in EPCs namely Diels-Alder catalysis and the Nazarov cyclization. In Figure 39 some of the EPC-catalyzed C-C bond formations are shown to illustrate the scope of these new phosphonium cation catalysts.

Figure 39 EPC-catalyzed carbon-carbon bond forming reactions and representative catalyst.87

A recent paper from Stephan et al. showed that EPCs are capable of catalyzing Diels-Alder reactions involving cyclohexa-1,3-diene and α, β-unsaturated ketones and Nazarov cyclisations.87_{In Figure 40}

the performed Diels-Alder reaction catalyzed by EPC is shown. The authors mentioned the use of EPC

24 (Figure 39) emerged as the optimal choice due to its performance and its easy accessibility.

Catalyst 24 converted a 2:1 combination of 28 and 27 into cycloadduct 29 with a 34% conversion at room temperature for 24 h with a high endo:exo ratio. Successively increasing the amount of the diene and then the catalyst led to enhanced conversion, reaching full conversion within 3 h with 8.0