Lewis Acidity Scale of Diaryliodonium Ions toward Oxygen, Nitrogen, and Halogen Lewis Bases

Article abstract of DOI:10.1021/jacs.9b12998

Equilibrium constants for the associations of 17 diaryliodonium salts Ar₂I⁺X^- with 11 different Lewis bases (halide ions, carboxylates, p-nitrophenolate, amines, and tris(p-anisyl)phosphine) have been investigated by titrations followed by photometric or conductometric methods as well as by isothermal titration calorimetry (ITC) in acetonitrile at 20 °C. The resulting set of equilibrium constants K_I covers 6 orders of magnitude and can be expressed by the linear free-energy relationship lg K_I = s_I LA_I + LB_I, which characterizes iodonium ions by the Lewis acidity parameter LA_I, as well as the iodonium-specific affinities of Lewis bases by the Lewis basicity parameter LB_I and the susceptibility s_I. Least squares minimization with the definition LA_I = 0 for Ph₂I⁺ and s_I = 1.00 for the benzoate ion provides Lewis acidities LA_I for 17 iodonium ions and Lewis basicities LB_I and s_I for 10 Lewis bases. The lack of a general correlation between the Lewis basicities LB_I (with respect to Ar₂I⁺) and LB (with respect to Ar₂CH⁺) indicates that different factors control the thermodynamics of Lewis adduct formation for iodonium ions and carbenium ions. Analysis of temperature-dependent equilibrium measurements as well as ITC experiments reveal a large entropic contribution to the observed Gibbs reaction energies for the Lewis adduct formations from iodonium ions and Lewis bases originating from solvation effects. The kinetics of the benzoate transfer from the bis(4-dimethylamino)-substituted benzhydryl benzoate Ar₂CH-OBz to the phenyl(perfluorophenyl)iodonium ion was found to follow a first-order rate law. The first-order rate constant k_obs was not affected by the concentration of Ph(C₆F₅)I⁺ indicating that the benzoate release from Ar₂CH-OBz proceeds via an unassisted S_N1-type mechanism followed by interception of the released benzoate ions by Ph(C₆F₅)I⁺ ions.

Full text of DOI:10.1021/jacs.9b12998

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Article
Lewis Acidity Scale of Diaryliodonium Ions toward Oxygen,
Nitrogen, and Halogen Lewis Bases
Robert J. Mayer,* Armin R. Ofial, Herbert Mayr, and Claude Y. Legault*
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ABSTRACT: Equilibrium constants for the associations of 17
diaryliodonium salts Ar2I+X− with 11 different Lewis bases (halide
ions, carboxylates, p-nitrophenolate, amines, and tris(p-anisyl)-
phosphine) have been investigated by titrations followed by
photometric or conductometric methods as well as by isothermal
titration calorimetry (ITC) in acetonitrile at 20 °C. The resulting set
of equilibrium constants KI covers 6 orders of magnitude and can be
expressed by the linear free-energy relationship lg KI = sI LAI + LBI,
which characterizes iodonium ions by the Lewis acidity parameter
LAI, as well as the iodonium-specific affinities of Lewis bases by the
Lewis basicity parameter LBI and the susceptibility sI. Least squares minimization with the definition LAI = 0 for Ph2I+ and sI = 1.00
for the benzoate ion provides Lewis acidities LAI for 17 iodonium ions and Lewis basicities LBI and sI for 10 Lewis bases. The lack of
a general correlation between the Lewis basicities LBI (with respect to Ar2I+) and LB (with respect to Ar2CH+) indicates that
different factors control the thermodynamics of Lewis adduct formation for iodonium ions and carbenium ions. Analysis of
temperature-dependent equilibrium measurements as well as ITC experiments reveal a large entropic contribution to the observed
Gibbs reaction energies for the Lewis adduct formations from iodonium ions and Lewis bases originating from solvation effects. The
kinetics of the benzoate transfer from the bis(4-dimethylamino)-substituted benzhydryl benzoate Ar2CH−OBz to the
phenyl(perfluorophenyl)iodonium ion was found to follow a first-order rate law. The first-order rate constant kobs was not affected
by the concentration of Ph(C6F5)I+ indicating that the benzoate release from Ar2CH−OBz proceeds via an unassisted SN1-type
mechanism followed by interception of the released benzoate ions by Ph(C6F5)I+ ions.
for catalyzing Diels−Alder reactions.6 Toste and co-workers
reported on the stereoselective C−H functionalization of cyclic
ethers with Lewis adducts of diaryliodonium salts with chiral
phosphates as key intermediates.7
INTRODUCTION
■
Diaryliodonium ions 1 (Scheme 1) are frequently employed
reagents for organic transformations.1 In particular, they are
Reports on the quantification of the Lewis acidities of
diaryliodonium ions are scarce. In 2003, Ochiai and co-workers
determined the binding constants of a small set of diaryl-
iodonium salts with 18-crown-6 by 1H NMR spectroscopy.8 In
2006, the same group reported on the binding constants of
Scheme 1. Structures of Diaryliodonium (1) and Iodolium
Ions (2)
9
nitrogen heteroarenes with Ph2I+BF4 . Okuyama studied
analogous reactions of cyclohexenyl phenyl iodonium ions with
bromide, and in this context also determined equilibrium
constants for the bromide addition.10 Recently, the Gutmann−
Beckett method11 was used to derive the relative Lewis
acidities of a small set of diaryliodonium salts from 31P NMR
chemical shift differences between triethylphosphine oxide and
−
useful for arylations, in which their ability to transfer an aryl-
group to oxygen, nitrogen, sulfur, halide, or carbon
nucleophiles, either with or without metal catalysis, gives
them a broad synthetic scope.2 Since the preassociation of the
nucleophile to the iodonium center is essential for the reaction
to proceed,3 the reactivity of diaryliodonium ions is largely
influenced by their Lewis acidity.
Received: December 7, 2019
More recently, applications exploiting the Lewis acidity of
iodonium ions as catalysts have also emerged. Liu, Han, and
co-workers employed diaryliodonium salts as Lewis acidic
catalysts for Mannich reactions,4 and the Huber group
reported the application of dibenzo[b,d]iodolium salts (2)5
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the corresponding Lewis adducts.12 This experimental Lewis
acidity scale was then combined with quantum chemical
calculations to extrapolate Lewis acidities of further iodine(III)
species.
Previously, Mayr and co-workers studied the equilibria for
the reactions of benzhydrylium ions with various C-, N-, O-,
P-, and S-centered Lewis bases in dichloromethane and
acetonitrile at 20 °C (eq 1).
RESULTS AND DISCUSSION
■
Structure of Iodonium Base Adducts. Solid-state (X-
ray) structures15 of the adducts of the diphenyliodonium ion
Ph2I+ (1a) with chloride,16 bromide,16 iodide,16 different
pyridines,9 18-crown-6,8 and DABCO17 have been reported in
the literature. The Lewis adducts of 1a with benzoate and p-
nitrophenolates (4a and 4b), were prepared in this work by
anion metathesis reactions (Scheme 2),18 which were driven
Scheme 2. Adducts of 1a with Benzoate Ion 3a and
Phenolate Ion 3e
The obtained equilibrium constants were correlated by eq 2:13
lg K = LA + LB
(2)
in which LA corresponds to a Lewis base independent Lewis
acidity parameter [LA = 0 for (4-MeOC6H4)2CH+] and LB is a
Lewis base parameter, which is assumed to generally hold for
reactions of Lewis bases with benzhydrylium ions.
We now report on the construction of a quantitative Lewis
acidity scale14 for diaryliodonium and iodolium ions by
measuring equilibrium constants KI for their association with
Lewis bases:
by the low solubility of potassium chloride in acetonitrile.
Sonication of a suspension of Ph2I+Cl− and potassium
benzoate (3a-K) or potassium 4-nitrophenolate (3e-K) in
acetonitrile led to precipitation of KCl and afforded
acetonitrile solutions of adducts 4a and 4b.
The Lewis adducts 4a and 4b crystallized upon partial
evaporation of the solvent. The iodonium benzoate 4a was
isolated in 53% yield. Crystals of 4a suitable for X-ray
crystallography were then obtained by diffusion of diethyl ether
into a concentrated acetonitrile solution of 4a. Iodonium p-
nitrophenolate 4b (isolated in 41% yield) was crystallized from
a concentrated acetonitrile solution.
Chart 1 lists the 17 Lewis acidic diaryliodonium (1a−n) and
dibenzo[b,d]iodolium ions (2a−c) covered in this work. Chart
2 gathers the structures of the 11 anionic and neutral entities
that were selected as reference Lewis bases.
In the solid-state structures of both adducts (Figure 1),
iodine adopts a square-planar coordination sphere with Ph−I−
Ph angles close to 90°. In 4a (benzoate adduct) this
coordination at iodine is achieved by formation of 2:2 adducts
where iodonium ions are linked by two benzoate ions in an
eight-membered ring. Analogous 2:2 adducts can be found in
the solid-state structure of the acetate salt of 1n.19 In crystals of
4b, strands are formed with the 4-nitrophenolate oxygen acting
as a μ2-bridging ligand.20 Both isolated adducts 4a and 4b are
stable at room temperature but undergo reductive elimination
to the corresponding ester/ether at higher temperatures, as
previously reported by Olofsson.21
Chart 1. Diaryliodonium and Iodolium Ions Investigated in
This Study
Determination of Equilibrium Constants. Equilibrium
constants for the formation of Lewis adducts from iodonium
species 1 (or 2) and Lewis bases 3 were determined by
different methods: conductometry, isothermal titration calo-
rimetry, UV/vis spectrometry, and NMR spectroscopic
titration.
Conductivity Measurements. Previous crystal structure
analyses8,9,16,17,19 as well as the results in this work (Figure 1)
illustrate the high tendency of the Lewis adducts of iodonium
ions to fill the coordination sphere of the iodine center by
formation of dimers or oligomers (Scheme 3). On the basis of
dimeric adducts found in the gas-phase in LC-MS experiments
and quantum-chemical calculations, Lee and co-workers
postulated that diaryliodonium chlorides also form 2:2 adducts
in solution.22
Chart 2. Reference Lewis Bases Used
To deepen our insight into the association equilibria of
iodonium salts, we used conductometry to study the degree of
dissociation of ammonium and iodonium salts in up to 1 mM
acetonitrile solutions.23 Figure 2 shows that the conductivities
of Ph2I+TfO− (1a), Bu4N+Cl−, and Bu4N+TfO− solutions
increase linearly with the salt concentrations in MeCN,
indicating almost complete dissociation of these salts and
B
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given by eq 4. The concentration of free ions can be calculated
when eq 4 is converted into eq 6 by using eq 5.
Fitting of [Ph2I+] calculated by eq 6 to the experimental
[Ph2I+] (blue dots in Figure 3) yields an excellent agreement of
Figure 1. X-ray structures of (A) diphenyliodonium benzoate 4a
(CCDC 1960557) and (B) diphenyliodonium 4-nitrophenolate 4b
(CCDC 1960556) at 173 and 143 K, respectively. Thermal ellipsoids
of 4a and 4b are depicted at a 50% probability level. For 4b, the
closest neighboring O and I atoms are shown (connected via dashed
lines). For 4b, the repetitive unit of the adduct strands is depicted
with the closest neighboring O and I atoms connected via dashed
lines. CCDC 1960556 and 1960557 contain the supplementary
crystallographic data for this paper. These data can be obtained free of
charge from The Cambridge Crystallographic Data Centre via
https://www.ccdc.cam.ac.uk/structures/.
Scheme 3. Possible Association Equilibria of Ph2I+ Ions with
Chloride Ions
Figure 3. Concentration of free Ph2I+ ions in a solution of Ph2I+Cl−
and fit (black line) of the experimental data with eq 6 (MeCN, 20
°C).
experimental and calculated ion concentrations for KI = (8.02
0.10) × 104 M−1 (black line in Figure 3). It is thus shown
that a 1:1 association model is sufficient to describe the
behavior of Ph2ICl in acetonitrile solution. Furthermore,
diffusion NMR experiments (DOSY) of the diphenyliodonium
ion in the presence of different counterions (Cl−, TfO−) (see
the Figure S3 for further details) indicated that the formation
of higher aggregates (dimers or oligomers) of Ph2ICl is
insignificant under our experimental conditions (<1 mM
solutions).
Isothermal Titration Calorimetry (ITC). In a micro-
calorimeter, small amounts of Lewis base (typically 40 × 6 μL
portions of a 10 mM solution) were added to a solution of the
iodonium salt (1.8 mL, 1 mM) at 20 °C, which gave rise to
heat evolution within the sample cell (Figure 4A). Integration
gives the heat evolution per injection. A plot of the heats of
reaction for the individual injections versus the molar ratio of
Lewis base over the iodonium ion yields a binding isotherm
(Figure 4B), which was analyzed with a 1:1 binding model
using the Affinimeter software package to afford the
equilibrium constant KI and the enthalpy of the binding
reaction ΔrH.26 The inflection point of the titration curve
provides experimental evidence for the 1:1 stoichiometry of
the interaction of Lewis bases and iodonium ions (shown for
all combinations in the Supporting Information).
Figure 2. Concentration-dependent conductivity of Ph2I+TfO− (1a),
Bu4N+TfO−, and Bu4N+Cl−, and Ph2I+Cl− (MeCN, 20 °C).
comparable molar conductivities of the involved ions in highly
diluted solution.24
At higher concentrations of Ph2I+TfO− 1a, the conductivity
increase loses its linear correlation with salt concentration due
to the formation of ion pairs (Figure S1). If compared to the
other ionic species in Figure 2, then the much lower
conductivities of the Ph2I+Cl− solutions in acetonitrile imply
that even at low concentration this salt is only partially
dissociated into free Ph2I+ and Cl− ions.25
Solvent Effects on Equilibrium Constants. NMR
From the measured conductivities of solutions of Ph2I+Cl−
and the specific conductivities of the fully dissociated salts
(Figure 2) at high dilution (Kohlrausch’s law), one can
calculate the concentration of the free ions [Ph2I+] (= [Cl−])
in Ph2ICl solutions in acetonitrile.
titrations were previously used to compare association
−
constants of Ph2I+PF6 (1b) with triethylphosphine oxide in
acetonitrile and dichloromethane at 25 °C, which showed that
KI(CH2Cl2)/KI(MeCN) = 3.8.12
Our ITC studies on the association of 1a with benzoate (3a)
showed effects in the same order of magnitude when
comparing acetonitrile with the less polar solvents dichloro-
methane [KI(CH2Cl2)/KI(MeCN) = 8.2] and THF
Assuming that Ph2ICl exists as a monomer (λ3-iodane) in
acetonitrile solution and not as a dimer or oligomer (i.e., Kdim
of Scheme 3 is negligibly small), the association constant KI is
C
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Figure 5. 19F DOSY NMR spectra (25 °C) of 1m in CD3CN (blue,
12 mM), CD2Cl2 (green, 11 mM), and THF-d8 (red, 8 mM). The 19
NMR spectrum on top refers to THF-d8.
F
Photometric Titrations (PT). KI (eq 3) for the association
of the colored Lewis base p-nitrophenolate (3e) with
diaryliodonium ions can be determined by direct spectropho-
tometric titration (Scheme 4). Stepwise addition of a Ph2I+
Figure 4. (A) Heat flow during the titration of Ph2I+TfO− (1a, 0.827
+
mM in MeCN) with NBu4 Cl− (10.0 mM in MeCN) at 20 °C (ITC
TfO− (1a) solution in acetonitrile to a solution of 3e (λmax
=
trace after baseline correction). (B) Plot of the integrated heats per
shot vs the equivalents of added chloride ions (black dots). The blue
line depicts the fitted titration curve (the first injection was omitted
from the correlation) giving KI for this individual experiment.
430 nm) in the same solvent led to a hypsochromic shift
(Figure 6A) of the absorption maximum, which we assigned to
the formation of the phenolate-iodonium complex.
Due to the overlap of the absorption bands of 3e and those
of corresponding iodonium adduct 4b, we determined the
molar fraction of phenolate χLB from Aλ at every step of the
titration (eq 12 in Scheme 4). As exemplified in Figure 6C, the
equilibrium constant KI was then obtained according to eq 8
(Scheme 4) from the slope of the linear correlation of the
[4b]/[3e] ratio (eq 9, Scheme 4) with the concentration of
[1a] (eq 11, Scheme 4). As full saturation cannot be achieved
[KI(THF)/KI(MeCN) = 20] (Table 1). The three equilibrium
constants (log KI) for the Lewis adduct formation 1a + 3a
27
correlate linearly with Reichardt’s ET values (r2 = 0.9867).
N
Table 1. Equilibrium Constants for the Lewis Adduct
Formation between Ph2I+TfO− (1a) and
Tetrabutylammonium Benzoate (3a) in Different Solvents
a
experimentally for some species, we used the value of Aλ in
ad
solvent
KI (M−1
)
the fitting process as a free variable to achieve an intercept of
the correlation in eq 8 (Scheme 4) that equals 0.
MeCN
CH2Cl2
THF
(1.76 0.16) × 105
(1.45 0.44) × 106
(3.58 0.16) × 106
The PT method is reliable for the determination of
equilibrium constants K < 106 M−1. Photometric titration of
p-nitrophenolate (3e) allowed to determine equilibrium
constants KI for the association of 3e with most of the
iodonium (1) ions investigated in this work (Supporting
Information).
Benzhydrylium-Indicator Method of Titration (BIMT).
The direct determination of the equilibrium constants by the
methods discussed so far are limited by physicochemical
properties of the Lewis base (charge, color), or the range, in
which reliable equilibrium constants can be obtained. We were,
therefore, seeking a more general method for the determi-
nation of KI.
Given the known reversibility of the reactions of many Lewis
bases with benzhydrylium ions (Ar′2CH+) in acetonitrile, we
chose colored benzhydrylium ions B1 and B229 as indicators to
establish equilibrium constants KI of diaryliodonium-Lewis
base reactions (Chart 3).
The BIMT relies on the thermodynamic relation illustrated
in Figure 7A. In analogy to the PT method and as shown in
Figure 7C, equilibrium constants KI are obtained from the
slope of the linear correlation of the [Ar2I−LB+]/[LB] ratio
a
ITC method, 20 °C.
We demonstrated above (Figure 2) that Ph2I+TfO− (1a) is
mostly dissociated in diluted acetonitrile solutions at <1 mM
and association of less than 5−15% is observed at
concentrations <2 mM (Figures S1 and S2). A different
situation is, however, observed in dichloromethane and THF.
In these less polar solvents, conductivity studies show that
Ph2I+TfO− (1a) exists mainly as ion pairs (Figure S5)28 whose
precise nature remains to be clarified. Additional evidence for
the influence of organic solvents on the extent of dissociation
of diaryliodonium triflates was obtained by 19F DOSY NMR
spectroscopy of iodonium salt 1m (Figure 5). In acetonitrile,
separate diffusion of the cation and the anion is detected. The
identical diffusion rates for both ions of 1m, however, provide
additional evidence for the presence of adducts (ion pairs) in
dichloromethane and THF. We have, therefore, focused our
further investigations on equilibria in acetonitrile solutions.
D
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Scheme 4. Determination of the Equilibrium Constants KI by Photometric Methods: Direct Titration with Colored Lewis
Bases (PT Method) and Benzhydrylium-Indicator Method of Titration (BIMT)
Chart 3. Benzhydrylium Ions Used as Indicators for the
BIMT
eq 2. Chart 3 lists the known Lewis acidity parameters LA for
B1 and B2,13 and Table 2 contains LB parameters for the
relevant charged and neutral Lewis bases (in acetonitrile).
Figure 7B illustrates how the titration of the benzhydrylium
ion B2 with BzO− (3a) and iodonium ion 1l can be monitored
by UV/vis spectrometry. A solution of blue benzhydrylium salt
B2-BF4 (stage A in Figure 7B) is first treated with Bu4N+BzO−
to convert approximately 50−70% of B2 into covalent,
colorless benzhydryl benzoate B2-OBz (step B in Figure
7B). Iodonium ion 1l (with the noncoordinating TfO−
counterion) was then added to the solution (step C in Figure
7B). The reaction of the free Lewis base (BzO−) with the
iodonium ion disturbs the benzhydrylium-Lewis base equili-
Figure 6. (A) Hypsochromic shift during the titration of 1a with 3e
+
brium (K in Figure 7A) and leads to regeneration of colored
benzhydrylium ion B2 from the reservoir of covalent B2-OBz.
The most striking feature of BIMT is that it is capable of
characterizing equilibrium constants involving a broad range of
colorless Lewis acids: Equilibrium constants KI from 102 to 108
M−1 can be determined, which is a significantly wider range
than with the PT method. By changing the reference
benzhydrylium ion (i.e., variation of K), this range can even
be extended. However, the equilibrium constant for the
(in MeCN at 20 °C: NBu4 ). (B) Plot of the molar fraction χLB vs
[1a]/[3e] and (C) determination of the equilibrium constant KI from
the slope of a linear plot of [4b]/[3e] vs [1a].
with [Ar2I+] (eq 8, Scheme 4). However, applying BIMT
requires additional knowledge of the equilibrium constants K
(eq 1) for the reactions of the Lewis bases with benzhydrylium
ions. Alternatively, K can be calculated from Lewis acidity and
Lewis basicity parameters LA and LB, respectively, by applying
E
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of benzhydrylium ions with Cl− (3i) and Br− (3j) have a high
uncertainty, and K remains to be determined for I− (3k).13
Comparison of Methods. To validate the consistency of
equilibrium constants KI derived by different methods, we
compared KI for the reactions of diphenyliodonium triflate
(1a) with the Lewis bases benzoate (3a), p-nitrophenolate
(3e), 4-(pyrrolidino)pyridine (3f), and chloride (3i). As
shown in Table 3, the agreement between KI from BIMT,
Table 3. Equilibrium Constants KI for the Adduct
Formation of Ph2I+ (1a) with Lewis Bases 3a, 3e, 3f, and 3i
a
Determined by Four Different Methods
Ph2I+ (1a) + LB
1a + 3a
method
KI (M−1
)
BIMT
ITC
BIMT
PT
(1.97 0.10) × 105
(1.76 0.16) × 105
(1.66 0.56) × 104
(8.02 0.33) × 103
(7.24 0.26) × 103
(1.06 0.05) × 102
(3.34 0.26) × 101,
(8.02 0.10) × 104
(6.29 0.42) × 104
1a + 3e
ITC
1a + 3f
BIMT
NMR
conduct.
ITC
b
1a + Cl− (3i)
a
b
MeCN, 20 °C. In CD3CN.
ITC, and PT methods is usually within the precision of the
measurements. Depending on the Lewis base/acid combina-
tion, the biggest deviations are found for BIMT applications:
While the errors are small when the equilibrium constant K for
the interactions of the Lewis base with the benzhydrylium ions
was directly measured, higher deviations (of factors up to 3)
have to be tolerated when K was extrapolated by eq 2 using
averaged LA and LB parameters because errors in K (eq 1)
propagate into KI (eq 14 in Scheme 4).
Figure 7. (A) Competing equilibria used for the determination of
equilibrium constants KI. (B) Representative BIMT experiment: A
solution of benzhydrylium tetrafluoroborate B2 (1.30 × 10−5 M, stage
A) was treated with Bu4N+BzO− (0.6 equiv), which generated a
mixture of B2 and the adduct B2-OBz (step B). Then a solution of 1l
(6.53 mM, counterion: TfO−) was added stepwise (step C), which
caused the incremental increase of the benzhydrylium ion
concentration [B2]. (C) Evaluation of KI according to eq 8 (Scheme
4) from the slope of the linear correlation of [1l-OBz]/[BzO−] vs
[1l].
+
The equilibrium constants for the reaction of NBu4 Cl− with
Ph2I+OTf− (1a) in acetonitrile determined by ITC agree well
with KI derived from the conductivity experiments (Figure 3),
which supports the assumed association model, that is, the
formation of a 1:1 Lewis acid−base adduct in solution.
However, differences in the total ion concentrations may
explain the 22% difference of KI for the 1a + 3i adduct
formation derived from conductometric titrations and ITC.
Solutions for conductometric measurements only contain
Ph2I+ and Cl− ions. In contrast, applying ITC requires the
presence of additional counterions, that is, TfO− (for 1a) and
Bu4N+ (for 3i), which may cause a change in the degree of ion
pairing that is seen in a slightly lower driving force for Lewis
adduct formation.
A similar trend is also found when comparing KI for 1a + 3f
obtained by applying the BIMT method with KI for the same
reaction from NMR titrations (Table 3).31 The equilibrium
constant KI under the conditions of the NMR spectroscopic
titration, which is by a factor of 3 lower than KI from BIMT
measurements, may indicate a higher degree of ion pairing (cf.
ref 23 and Figures S1 and S2) owing to the more than 5-fold
higher total ion concentrations in NMR experiments than in
photometric or ITC experiments.
a
Table 2. Lewis Basicities LB
Lewis base
LB (in MeCN)
ref
b
b
b
b
3a
3b
3c
3d
3e
3f
3g
3h
3i
17.54
16.66
15.64
15.28
16.98
17.39
15.48
17.26
≈11
this work
this work
this work
this work
30
13
13
13
13
3j
≈8
13
3k
not known
a
b
As defined in eq 2. LB was determined by following the method
described in ref 13. For details, see the Supporting Information.
Table 4 lists all equilibrium constants KI for reactions of
Lewis bases with iodonium (1) and iodolium ions (2) in
acetonitrile (at 20 °C) that were determined in this work.
Almost identical equilibrium constants KI were measured for
the adduct formation of phenolate ion 3e with the
diphenyliodonium triflate (1a) and the diphenyliodonium
reaction of the benzhydrylium ion with the Lewis base K (eq
1) needs to be known with sufficient precision. As a
consequence, BIMT was inapplicable for determining KI of
halide anions 3i−k: The binding constants K for the reaction
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a
Table 4. Equilibrium Constants KI (M−1) for the Reactions of Diaryliodonium Ions 1 and 2 with Lewis Bases 5
b
c
Lewis acid
LAI
Lewis base
LBI
sI
KIexptl
method
KIcalcd
KIcalcd/KIexptl
d
1a
1a
1a
1a
1a
1a
1a
1a
1a
1a
1b
1c
1c
1c
1d
1d
1d
1d
1d
1e
1e
1e
1e
1f
0.00
0.00
0.00
0.00
0.00
0.00
0.00
0.00
0.00
0.00
−0.07
0.60
0.60
0.60
1.27
1.27
1.27
1.27
1.27
3.21
3.21
3.21
3.21
0.89
0.89
0.89
0.89
0.89
0.89
0.89
0.89
0.89
0.89
−0.26
−0.26
0.53
0.53
0.53
−0.76
−0.76
−0.76
−0.76
−0.76
−0.11
−0.11
−0.11
−0.24
−0.24
0.96
3a
3b
3c
3d
3e
3f
3g
3i
3j
3k
3e
3a
3e
3f
3a
3e
3f
3g
3i
3a
3e
3f
3g
3a
3b
3c
3d
3e
3f
3g
3i
3j
3k
3a
3e
3a
3e
3f
3a
3e
3i
5.33
4.78
4.28
4.52
3.96
2.21
2.25
4.71
4.29
3.62
3.96
5.33
3.96
2.21
5.33
3.96
2.21
2.25
4.71
5.33
3.96
2.21
2.25
5.33
4.78
4.28
4.52
3.96
2.21
2.25
4.71
4.29
3.62
5.33
3.96
5.33
3.96
2.21
5.33
3.96
4.71
4.29
3.62
5.33
3.96
2.21
5.33
3.96
5.33
3.96
2.21
5.33
3.96
2.21
5.33
3.96
4.71
5.33
4.78
4.28
4.52
1.00
0.94
0.92
0.88
0.91
0.31
0.85
0.89
0.84
0.80
0.91
1.00
0.91
0.31
1.00
0.91
0.31
0.85
0.89
1.00
0.91
0.31
0.85
1.00
0.94
0.92
0.88
0.91
0.31
0.85
0.89
0.84
0.80
1.00
0.91
1.00
0.91
0.31
1.00
0.91
0.89
0.84
0.80
1.00
0.91
0.31
1.00
0.91
1.00
0.91
0.31
1.00
0.91
0.31
1.00
0.91
0.89
1.00
0.94
0.92
0.88
(1.97 0.10) × 105
(4.95 0.45) × 104
(1.68 0.09) × 104
(2.94 0.06) × 104
(8.02 0.33) × 103
(1.06 0.05) × 102
(2.66 0.23) × 102
(6.29 0.42) × 104
(2.32 0.06) × 104
(4.39 0.44) × 103
(7.78 0.65) × 103
(9.52 0.10) × 105
(3.18 0.12) × 104
(1.85 0.15) × 102
(5.10 0.09) × 106
(1.37 0.08) × 105
(8.71 0.35) × 102
(1.32 0.02) × 103
(5.95 0.14) × 105
(4.07 0.11) × 108
(6.57 0.58) × 106
(8.68 0.86) × 102
(1.18 0.11) × 105
(1.76 0.03) × 106
(5.83 1.08) × 105
(1.59 0.09) × 105
(2.48 0.06) × 105
(2.14 0.28) × 104
(2.92 0.20) × 102
(8.58 0.14) × 102
(4.23 0.19) × 105
(1.03 0.05) × 105
(2.21 0.07) × 104
(1.27 0.02) × 105
(4.85 0.14) × 103
(1.08 0.03) × 106
(2.01 0.13) × 104
(1.68 0.17) × 102
(3.47 0.05) × 104
(2.20 0.48) × 103
(1.16 0.06) × 104
(4.04 0.22) × 103
(9.71 2.12) × 102
(2.55 0.34) × 105
(4.62 0.11) × 103
(1.27 0.05) × 102
(8.53 0.77) × 104
(8.28 0.05) × 103
(2.97 0.02) × 106
(4.36 0.07) × 104
(3.19 0.17) × 102
(3.03 0.12) × 106
(2.35 0.03) × 104
(6.62 0.09) × 102
(5.92 0.98) × 104
(8.60 1.14) × 103
(1.68 0.03) × 104
(1.27 0.05) × 107
(3.12 0.28) × 106
(9.69 0.74) × 105
(1.41 0.09) × 106
BIMT
2.13 × 105
5.98 × 104
1.91 × 104
3.30 × 104
9.02 × 103
1.62 × 102
1.78 × 102
5.12 × 104
1.94 × 104
4.14 × 103
−
1.1
1.2
1.1
1.1
1.1
1.5
0.67
0.81
0.84
0.94
−
0.90
1.0
1.4
0.78
0.95
0.47
1.6
1.2
0.86
1.2
d
BIMT
d
BIMT
d
BIMT
PT
e
BIMT
d
BIMT
ITC
ITC
ITC
PT
d
BIMT
8.58 × 105
3.21 × 104
2.51 × 102
3.99 × 106
1.31 × 105
4.07 × 102
2.12 × 103
6.85 × 105
3.50 × 108
7.72 × 106
1.65 × 103
9.36 × 104
1.65 × 106
4.09 × 105
1.25 × 105
1.99 × 105
5.83 × 104
3.08 × 102
1.00 × 103
3.13 × 105
1.07 × 105
2.14 × 104
1.18 × 105
5.26 × 103
7.26 × 105
2.76 × 104
2.38 × 102
3.74 × 104
1.84 × 103
1.10 × 104
4.52 × 103
1.02 × 103
1.64 × 105
7.09 × 103
1.50 × 102
1.24 × 105
5.50 × 103
1.96 × 106
6.82 × 104
3.26 × 102
1.67 × 106
5.88 × 104
3.09 × 102
8.81 × 104
4.03 × 103
2.34 × 104
1.71 × 107
3.68 × 106
1.08 × 106
1.56 × 106
PT
e
BIMT
d
BIMT
PT
e
BIMT
d
BIMT
ITC
d
BIMT
d
BIMT
e
BIMT
1.9
d
BIMT
0.79
0.94
0.70
0.79
0.80
2.7
d
BIMT
d
1f
1f
1f
1f
1f
1f
1f
1f
BIMT
d
BIMT
d
BIMT
PT
e
BIMT
1.1
1.2
d
BIMT
ITC
ITC
ITC
0.74
1.0
0.97
0.93
1.1
0.67
1.4
1.4
1f
d
1g
1g
1h
1h
1h
1i
1i
1i
1i
1i
1j
1j
1j
1k
1k
1l
1l
1l
1m
1m
1m
1n
1n
1n
2a
2a
2a
2a
BIMT
PT
d
BIMT
PT
e
BIMT
d
BIMT
1.1
PT
0.84
0.94
1.1
1.1
0.64
1.5
1.2
1.5
0.66
0.66
1.6
1.0
0.55
2.5
0.47
1.5
0.47
1.4
1.4
1.2
1.1
1.1
ITC
ITC
ITC
3j
3k
3a
3e
3f
3a
3e
3a
3e
3f
3a
3e
3f
3a
3e
3i
d
BIMT
PT
e
BIMT
d
BIMT
PT
d
BIMT
0.96
0.96
0.89
0.89
PT
e
BIMT
d
BIMT
PT
e
0.89
BIMT
d
−0.38
−0.38
−0.38
1.90
1.90
1.90
BIMT
PT
ITC
d
3a
3b
3c
3d
BIMT
d
BIMT
d
BIMT
d
1.90
BIMT
G
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J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
pubs.acs.org/JACS
Article
Table 4. continued
b
c
Lewis acid
LAI
Lewis base
LBI
sI
KIexptl
method
KIcalcd
KIcalcd/KIexptl
d
d
d
2a
2a
2a
2a
2a
2a
2b
2b
2c
2c
1.90
1.90
1.90
1.90
1.90
1.90
2.79
2.79
0.78
0.78
3e
3f
3g
3i
3.96
2.21
2.25
4.71
4.29
3.62
5.33
3.96
5.33
3.96
0.91
0.31
0.85
0.89
0.84
0.80
1.00
0.91
1.00
0.91
(9.87 0.38) × 105
(1.00 0.91) × 103
(9.89 0.80) × 102,
(2.27 0.14) × 106
(7.43 0.05) × 105
(1.35 0.10) × 105
(9.08 0.16) × 107,
(4.72 0.08) × 106,
(5.23 0.33) × 105,
(1.25 0.06) × 105,
BIMT
BIMT
BIMT
ITC
ITC
ITC
4.92 × 105
6.42 × 102
7.28 × 103
2.48 × 106
7.61 × 105
1.40 × 105
1.31 × 108
3.15 × 106
1.29 × 106
4.65 × 104
0.50
0.64
7.4
1.1
1.0
1.0
1.4
0.67
2.5
0.37
df
,
3j
3k
3a
3e
3a
3e
g
g
e
e
d
d
d
d
BIMT
BIMT
BIMT
BIMT
a
b
Determined in acetonitrile at 20 °C and compared with the equilibrium constants derived by applying eq 18. Averaged from at least three
measurements. Calculated by applying LAI, LBI, and sI in eq 18. With B2 as competition partner. With B1 as competition partner. Not used for
c
d
e
f
g
the least-squares minimization. Measurements of lower quality.
2
Δ = ∑ (lg KIexptl − lg KIcalcd 2
)
hexafluorophosphate (1b) indicating the marginal counterion
effect of these noncoordinating anions. Interactions of
iodonium ions with tris(p-anisyl)phosphine (3h) are so weak
that KI could not even be determined for the most Lewis acidic
iodonium ion 1e.32
Correlation Analysis of the Equilibrium Constants KI.
As shown in Figure 8, in general, the ordering of the Lewis
= ∑ (lg KIexptl − (sILAI + LBI))2
(19)
The resulting iodonium-specific Lewis basicity parameters
LBI (and sI) for 3 as well as the Lewis acidities LAI of
iodonium ions 1 and 2 are listed in the left columns of Table 4.
The columns on the right show that the equilibrium constants
KIcalcd calculated by eq 18 from sI, LAI, and LBI (obtained by
least-squares minimization of Δ2, eq 19) deviate from the
individual experimental KIexptl by less than a factor of 3
(exception: 2a + 3g). The quality of the linear correlations is
illustrated by Figure 8, which also indicates the different slopes
sI of the correlations. To assess the overall quality of our
correlation, Figure S7 shows the plot of calculated lg KIcalcd
versus experimental lg KIexptl values, giving an excellent linear
correlation (R2 = 0.987).
The experimental equilibrium constant KI for the adduct
formation between quinuclidine (3g) and iodolium ion 2a is
7.5 times smaller than KIcalcd. This observation might be caused
by repulsive interactions between the sterically demanding
quinuclidine (3g) in combination with the rigid structure of
iodolium ion 2a which enforces binding within its plane.
Structure−Reactivity Relationships of Iodonium Ions.
The Lewis acidity parameters for iodonium ions LAI, which
correspond to averaged lg KI values, can now be used to
elucidate the impact of steric and electronic effects on the
observed Lewis acidities (Scheme 5). The Lewis acidities LAI
of dibenzo[b,d]iodolium ions 2 are about 2 orders of
magnitude greater than those of the analogously substituted
diaryliodonium ions 1 (Scheme 5A).
Figure 8. Plot of lg KI for the reactions of diaryliodonium ions 1 (and
2) with Lewis bases 3 versus their Lewis acidity parameters LAI (all
individual correlations are shown in Figure S6).
Comparison of 1a with 1g and of 1f with 1h (Scheme 5B)
shows that replacement of phenyl by mesityl reduces the Lewis
acidity only slightly. The equilibrium constants KI for
iodonium ions bearing one mesityl moiety are by a factor of
2 smaller than for those that carry a phenyl group instead.
Exchanging one phenyl group in Ph2I+ by a 2,4,6-
trimethoxyphenyl (TMP) group (1a vs 1n) reduces the
Lewis acidity by 0.38 LAI units. The large atomic radius of
iodine (140 pm) and the fact that the plane of the aryl groups
is perpendicular to the C−I−C plane may explain the
observation that the steric bulk of the mesityl or the TMP
group has only a small effect on the Lewis acidity. This
rationalization is also in line with the solid-state structures of
the Lewis adducts in Figure 1, which illustrate that the two aryl
groups adopt a conformation perpendicular to the C−I−C
plane and that the distance between iodine and the
acidities of iodonium ions is the same with respect to different
Lewis bases, but the relative Lewis basicities slightly depend on
the nature of the iodonium ion. The equilibrium constants can
be expressed by the linear-free energy relationship (eq 18),
where the Lewis acidities of the diaryliodonium (and
iodolium) ions are characterized by the parameter LAI and
the Lewis basicities are characterized by the parameter LBI and
the Lewis base specific susceptibilities by sI.
lg KI = sI·LAI + LBI
(18)
These parameters were determined by least-squares mini-
mization of Δ2 (eq 19) from 70 equilibrium constants KI
(Table 4) with the definitions LAI(Ph2I+) = 0.00 and sI(BzO−)
= 1.00.
H
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Journal of the American Chemical Society
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Article
values were then analyzed according to the van ’t Hoff
equation to calculate ΔrH and ΔrS.
Scheme 5. Comparison of the Lewis Acidities LAI of
Different Iodonium Ions
For all studied systems, there is a fair agreement between
directly measured enthalpies (by ITC) with those based on a
van ’t Hoff analysis (KI from ITC, PT and BIMT). However,
van ’t Hoff derived enthalpies ΔrH are systematically by 2−5 kJ
mol−1 less negative than the directly measured heats of
reaction (ΔrH). From the entries in Table 5, it is evident that
there is a significant entropic contribution to the interactions
of Ph2I+ with Lewis bases. The positive entropy ΔrS for
reactions of 1a with anions can be rationalized by the higher
ordering of the solvent molecules around the ionic reactants.
Adduct formation reduces the electron density at the atoms of
both reaction centers resulting in less positive ΔrS for the
highly delocalized 4-nitrophenolate (3e): The importance of
stabilization by the solvation shell, and thus the entropic
component, is reduced compared to the benzoate ion (3a)
with the more localized negative charge.
In reactions of 1a with neutral Lewis base 3g the opposite
entropic effect is observed: ΔrS is now negative, causing a
reduction of the exergonicity for the formation of the
corresponding Lewis adduct. Notably, the reaction enthalpy
of association for 3g is by more than 30 kJ/mol more negative
than for the association with 3a though the latter has a higher
affinity towards iodonium ions (ΔrG).
The strong entropic effect is furthermore illustrated in a
comparison of the reactions of various iodonium ions with the
same Lewis base (e.g., chloride (3i)), which shows that the
increase of the binding constants is mostly due to an increase
of entropy, while the reaction enthalpy is almost constant
(Table 6).
Kinetics of the Iodonium-Induced Ionization of
Benzhydryl Benzoates. Rate constants for ionizations of
covalent benzhydryl benzoate B2-OBz in the presence of
variable amounts of the Lewis acidic iodonium ion 1e in
acetonitrile were determined using stopped-flow UV/vis
photometry by monitoring the increasing absorption of the
reaction mixture caused by the released benzhydrylium ion B2.
As depicted in Scheme 6, the heterolysis of B2-OBz may
proceed through two mechanistic pathways: (A) The
iodonium ion attacks at the benzhydryl benzoate to yield
complex C, which may be an intermediate or a transition state
on the way to the final products. In pathway (B), the
heterolytic cleavage of the benzhydryl benzoate occurs without
assistance by the iodonium ion, and the iodonium ion Ar2I+
just traps the benzoate ion generated in an SN1 process.
coordinating oxygen atom in the Lewis base is kept relatively
long.
Why does the exchange of phenyl by mesityl have a larger
effect on LAI (ΔLAI = −0.51) when 1g and 1i are compared?
According to calculations,12 the Lewis base preferentially binds
trans to the most electron-deficient aryl group when the T-
shaped adduct is formed (Scheme 5C). Since only in the
comparison of 1g/1i the aryl group is exchanged, which will be
located trans to the Lewis base in the adduct, the change from
phenyl to mesityl has the strongest impact. Nevertheless, even
in this case, the equilibrium constant KI is reduced only by a
factor of 3.
Reaction Enthalpies of Lewis Adduct Formation. The
ITC measurements yield not only the equilibrium constant KI
for a certain Lewis adduct formation but also the related
enthalpy of reaction (ΔrH) the combination of which yields
the entropy of reaction (ΔrS) (Table 5).
For examining the consistency of the thermodynamic data
obtained from ITC measurements at 20 °C, we additionally
determined temperature-dependent equilibrium constants KI
by ITC, PT, and BIMT.33 The temperature-dependent KI
a
Table 5. Thermodynamics of the Lewis Adduct Formation between Ph2I+TfO− (1a) and Lewis Bases 3
entry
Lewis base
method
ITC (20 °C)
ITC (20 to 50 °C)
ITC (20 °C)
ITC (20 to 40 °C)
PT (10 to 40 °C)
BIMT (−10 to 20 °C)
ITC (20 °C)
ITC (20 to 50 °C)
ITC (20 °C)
ITC (20 °C)
ΔrG (kJ mol−1
−29.4 0.2
)
ΔrH (kJ mol−1
)
ΤΔrS (kJ mol−1
)
ΔrS (J mol−1 K−1
)
1
2
3
4
BzO− (3a)
BzO− (3a)
−18.1 0.5
−13.2 1.6
−16.9 0.4
−14.3 2.0
−14.0 0.6
−50.6 3.5
−11.2 0.3
−9.7 0.8
−10.8 0.1
−9.8 0.3
11.4 0.3
16.3 1.6
5.0 0.4
7.5 1.9
7.9 0.6
−37.2 3.8
15.7 0.2
17.1 0.8
13.7 0.1
10.6 0.5
38.9 1.2
55.6 5.3
17.2 1.4
25.5 6.5
26.9 2.2
−127 13
53.6 0.4
58.3 2.6
56.6 0.5
36.2 1.8
b
b
p-nitrophenolate (3e)
p-nitrophenolate (3e)
p-nitrophenolate (3e)
quinuclidine (3g)
Cl− (3i)
−21.9 0.3
c
b
5
−21.9 0.1
−13.6 0.2
−26.9 0.1
b
6
7
8
9
b
Cl− (3i)
Br− (3j)
−24.5 0.1
−20.4 0.2
10
I− (3k)
a
b
In MeCN, 20 °C. Equilibrium constants KI were determined at variable temperatures and used in a van ’t Hoff plot to determine ΔrH and ΔrS
c
−
(see the Supporting Information). Ph2I+PF6 (1b) was used.
I
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Journal of the American Chemical Society
pubs.acs.org/JACS
Article
a
Table 6. Thermodynamics of the Lewis Adduct Formation between Chloride Ion 3i and Iodonium Salts 1a, 1d, 1f, 1i, and 2a
iodonium salt
KI (M−1
)
ΔrG (kJ mol−1
)
ΔrH (kJ mol−1
)
ΤΔrS (kJ mol−1
)
ΔrS (J mol−1 K−1
)
1i
1a
1f
1d
2a
(1.16 0.06) × 104
(6.29 0.42) × 104
(4.23 0.19) × 105
(5.95 0.14) × 105
(2.27 0.14) × 106
−22.8 0.1
−26.9 0.1
−31.6 0.1
−32.4 0.1
−35.7 0.1
−10.2 0.2
−11.2 0.3
−12.2 0.1
−10.5 0.1
−13.5 0.2
12.6 0.3
15.7 0.2
19.4 0.1
21.9 0.1
22.2 0.1
43.1 1.1
53.6 0.4
66.2 0.4
74.6 0.3
75.8 0.5
a
ITC method, MeCN, 20 °C.
The rate constant khet(20 °C) for the unassisted heterolysis of
the covalent B2-OBz in acetonitrile can independently be
estimated from the ratio khet = k2/K according to Scheme 7.
The equilibrium constant for the reaction of benzhydrylium
ion B2 with 3a in acetonitrile can be calculated to be K = 5.2 ×
107 M−1 by eq 2, from the Lewis acidity LA = −9.82 of B2 and
the Lewis basicity LB = 17.54 of BzO− (3a). The second-order
rate constant k2 = 6.2 × 106 M−1 s−1 for the reaction of B2 (E
= −7.02) with 3a (N = 16.45, sN = 0.72)34 in acetonitrile can
be calculated by using the Mayr−Patz equation (eq 20).35
Scheme 6. Pathways for the Benzoate Transfer from a
Benzhydrylium Benzoate to an Iodonium Ion
When the formation of benzhydrylium ion B2 from B2-OBz
(in the presence of 30% B2) was monitored photometrically
lg k2(20°C) = sN(N + E)
(20)
The ratio k2/K = 0.12 s−1 equals the unassisted heterolysis rate
constant khet. Obviously, the presence of Lewis acidic
iodonium ions 1e at a submillimolar concentration does not
significantly increase the rate of the heterolytic cleavage of the
Ar2CH−OCOPh bond in B2-OBz.
with different concentrations of added Ph(C6F5)I+PF6 (1e),
−
first-order kinetics were observed. The observed first-order rate
constants kobs were independent of the concentration of the
added iodonium ion (1e, Figure 9). From this observation, one
In line with this interpretation, Huber found almost identical
rates for the solvolysis reaction of Ph2CH−Cl in wet
acetonitrile with both iodolium ions 2a and 2b6 even though
their Lewis acidity differs by 1 order of magnitude according to
our measurements.
Comparison of Lewis Basicities toward Iodine- and
Carbon-Centered Lewis Acids. The blue correlation line in
Figure 10 shows that the equilibrium constants KI for the
reactions of p-substituted benzoate ions 3a−d with dipheny-
liodonium ion 1a correlate linearly with the Lewis basicities LB
(toward benzhydrylium ions) of these anions. The slope of this
correlation (0.41) implies that variation of the substituents
affects the equilibrium constants toward iodonium ions by 41%
of the effect on the equilibrium constants of reactions with
Figure 9. Time-dependent absorbance A (at 605 nm) for the reaction
of B2-OBz (≈ 6 × 10−6 M) with different concentrations of the
iodonium ion 1e (0.1 to 0.5 mM) in acetonitrile at 20 °C.
can already conclude that the iodonium ions 1e are not
involved in the rate-determining step. Benzoate ions are
generated by unassisted heterolytic cleavage of B2-OBz (path
B in Scheme 6), and iodonium ion 1e just serves as a trap for
the released benzoate ions.
This conclusion is further corroborated by the similarity of
the measured kobs (0.11 s−1, Figure 8) and the rate constant for
the unassisted heterolysis khet (0.12 s−1) derived in Scheme 7.
Scheme 7. Estimating the Ionization Constant khet of B2-
OBz in Acetonitrile from Reactivity and Lewis Acidity/
Basicity Parameters
Figure 10. Plot of lg KI for the reactions of Ph2I+ (1a) with various
Lewis bases (MeCN, 20 °C) versus the Lewis basicity LB of the same
set of Lewis bases in adduct formations with benzhydrylium ions
(with data from Table 2). Blue correlation: lg KI = 0.41LB − 2.0, R2 =
0.832.
J
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pubs.acs.org/JACS
Article
benzhydrylium ions. The weak interaction between iodonium
ions and these Lewis bases, which can be derived from this
small slope (e.g., for benzoates), is in line with the observation
that also the UV/vis spectrum of p-nitrophenolate is only
slightly changed during association with an iodonium ion
(Figure 6A).
The scatter of the other entries in Figure 10 shows that
Lewis basicity toward iodonium ions is controlled by other
factors than basicity toward carbon centers. While benzoate
(3a), 4-nitrophenolate (3e), and 4-pyrrolidinopyridine (3f)
have comparable Lewis basicities toward benzhydrylium ions,
their ability to form Lewis adducts with Ph2I+ (1a) differs by
more than 3 orders of magnitude. The equilibrium constant KI
for the reaction of Ph2I+ (1a) with tris(4-methoxyphenyl)-
phosphine (3h) was so small that we were not able to
determine it though its LB toward Ar2CH+ (17.26) is
comparable to that of 3a, 3e, and 3f. In contrast, bromide
(3j) and chloride ions (3i) have KI toward Ph2I+ similar to
those of benzoate and p-nitrophenolate, while their Lewis
basicities toward benzhydrylium ions differ by almost 9 orders
of magnitude.
Figure 11. Lewis acidities LAI of iodonium ions in acetonitrile.
Only for structurally closely related Lewis bases (p-
substituted benzoates) do the Lewis basicities toward
iodonium ions correlate with the Lewis basicities toward
carbenium ions. In contrast, the relative Lewis basicities of
halide anions, O-, N-, and P-centered Lewis bases, toward
iodonium ions differ dramatically from their relative Lewis
basicities toward benzhydrylium ions (Figure 10).
Kinetic investigations of the iodonium-ion-promoted ioniza-
tion of benzhydryl benzoate B2-OBz showed that not even the
most Lewis acidic iodonium ion in this work 1e accelerated the
ionization of B2-OBz. Iodonium ion rather quenches the
benzoate ions generated by unassisted ionization of B2-OBz
and thus increases the degree of ionization of the
benzhydrylium benzoate (Le Chatelier principle).
CONCLUSION
■
In previous work, it has been demonstrated that the
equilibrium constants for the reactions of Lewis bases with
benzhydrylium ions can be described by the sum of a base-
independent Lewis acidity parameter LA and an acid-
independent Lewis basicity parameter LB (eq 2), that is, the
relative strengths of different Lewis bases are the same toward
benzhydrylium ions of low and high Lewis acidity. The
situation is more complex for the formation of Lewis adducts
from iodonium ions and Lewis bases.
ASSOCIATED CONTENT
* Supporting Information
The Supporting Information is available free of charge at
https://pubs.acs.org/doi/10.1021/jacs.9b12998.
■
In organic solvents, diaryliodonium triflates and hexafluor-
phosphates can form ion pairs to a certain extent. While ion-
pairing is negligible in diluted acetonitrile solution (<2 mM,
Figures S1 and S2), in dichloromethane and THF it is already
significant in the submillimolar concentration range (Figure
S5).
Using conductometric and photometric titration (PT), ITC,
and titrations using colored benzhydrylium ions B1 and B2 as
indicators (BIMT), equilibrium constants for the coordination
of iodonium ions with Lewis bases (halide, benzoate, and
phenolate ions, as well as amines) have been determined in
acetonitrile at 20 °C.
By subjecting these 70 equilibrium constants KI to a least-
squares minimization on the basis of eq 18, Lewis acidity
parameters LAI for different types of diaryliodonium ions have
been obtained which apply for associations with all types of
Lewis bases (Figure 8). Figure 11 shows that the acidities of
the diaryliodonium ions increase with increasing electron-
accepting abilities of the substituents and cover 4 orders of
magnitude from the weakest Lewis acid, the dimesityliodonium
ion (1i), to the strongest one, phenyl(pentafluorophenyl)-
iodonium ion (1e). Cyclic iodonium ions 2a−c have a
considerably higher Lewis acidity than their acyclic analogs.
The Lewis acidity scale for diaryliodonium ions derived in this
work is in line with the relative acidities determined by the
Gutmann−Beckett method.12 While NMR experiments, which
are usually carried out at higher concentrations, give the same
ranking of Lewis acidities, the absolute values of KI can be
expected to be somewhat lower than the KI determined in this
study due to the effects of ion pairing.
sı
Synthetic procedures, details of the determination of
equilibrium constants K and KI, details of kinetic
measurements (PDF)
Crystallographic data for 4a and 4b (CIF, CIF)
AUTHOR INFORMATION
Corresponding Authors
■
Robert J. Mayer − Department Chemie, Ludwig-Maximilians-
̈
̈
̈
Universitat Munchen, 81377 Munchen, Germany;
orcid.org/0000-0002-9864-7042; Email: robert.mayer@
cup.uni-muenchen.de
Claude Y. Legault − University of Sherbrooke, Department of
Chemistry, Centre in Green Chemistry and Catalysis,
́
Sherbrooke, Quebec J1K 2R1, Canada; orcid.org/0000-
0002-0730-0263; Email: claude.legault@usherbrooke.ca
Authors
Armin R. Ofial − Department Chemie, Ludwig-Maximilians-
̈
̈
̈
Universitat Munchen, 81377 Munchen, Germany;
orcid.org/0000-0002-9600-2793
Herbert Mayr − Department Chemie, Ludwig-Maximilians-
̈
̈
̈
Universitat Munchen, 81377 Munchen, Germany;
orcid.org/0000-0003-0768-5199
Complete contact information is available at:
https://pubs.acs.org/10.1021/jacs.9b12998
K
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J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Author Contributions
The manuscript was written through contributions of all
authors. All authors have given approval to the final version of
the manuscript.
pubs.acs.org/JACS
Article
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Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
We thank Prof. Stefan Huber and Flemming Heinen (Ruhr-
̈
Universitat Bochum, Germany) for providing samples of the
iodolium salts 2a−c and for helpful discussions. We thank
Nathalie Hampel (LMU Munchen) for the experimental
̈
support and Dr. Peter Mayer (LMU Munchen) for single
̈
crystal X-ray structure determinations. This work was
supported by the National Science and Engineering Research
Council (NSERC) of Canada, the Canada Foundation for
́
Innovation (CFI), the Fonds de Recherche du Quebec −
Nature et Technologies (FRQNT), and the Universite de
Sherbrooke. R.J.M. thanks the Fonds der Chemischen
́
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Industrie for a Kekule fellowship. The authors want to thank
one anonymous referee for alerting us of the significance of ion
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