Characterization of species present in aqueous solutions of [hydroxy(mesyloxy)iodo]benzene and [hydroxy(tosyloxy)iodo]benzene

Article abstract of DOI:10.1021/ja971751c

Upon solution in water, both [hydroxy(mesyloxy)iodo]benzene and [hydroxy(tosyloxy)iodo]benzene undergo complete ionization to give the hydroxy(phenyl)iodonium ion (PhI⁺OH) and the corresponding sulfonate ion (RSO₂O^-) as fully solvated species, i.e., 'free' ions. The phenyliodonium solution species do not form ion pairs with the organosulfonate ions. The hydroxy(phenyl)iodonium ion is presumed to be ligated with at least one water molecule at an apical site of the iodine(m) atom originally occupied by the sulfonate ion. In view of the relative basicities of HO^- and H₂O, the hydroxy ligand of the [hydroxy(aquo)iodo]benzene ion (PhI⁺(OH₂)OH) is expected to be strongly bound and the water ligand is expected to be weakly bound to the iodine(m) center. This species has a pK(A) at (4.30 ± 0.05). PhI⁺(OH₂)OH and its conjugate base are present in equilibrium with the [hydroxy(auqo)]-μ-oxodiphenyldiiodine cation (Ph(HO)I-O-I⁺(OH₂)Ph). This μ-oxo dimer is present at significant levels even in relatively dilute solutions as the combination equilibrium constant is (540 ± 50). This dimer can be protonated, and the pK(A) of the conjugate acid is ~2.5. The equilibrium constant for dimerization of [oxo(aquo)iodo]benzene (PhI⁺(OH₂)O^-), the most important monomer in acidic solutions, is ~8.6.

Full text of DOI:10.1021/ja971751c

9614
J. Am. Chem. Soc. 1997, 119, 9614-9623
Characterization of Species Present in Aqueous Solutions of
[Hydroxy(mesyloxy)iodo]benzene and
[Hydroxy(tosyloxy)iodo]benzene
Helen Wilkinson Richter,* Brian R. Cherry, Teresa D. Zook, and Gerald F. Koser
Contribution from the Department of Chemistry, The UniVersity of Akron,
Akron, Ohio 44325-3601
ReceiVed May 28, 1997X
Abstract: Upon solution in water, both [hydroxy(mesyloxy)iodo]benzene and [hydroxy(tosyloxy)iodo]benzene undergo
complete ionization to give the hydroxy(phenyl)iodonium ion (PhI+OH) and the corresponding sulfonate ion (RSO2O-)
as fully solvated species, i.e., “free” ions. The phenyliodonium solution species do not form ion pairs with the
organosulfonate ions. The hydroxy(phenyl)iodonium ion is presumed to be ligated with at least one water molecule
at an apical site of the iodine(III) atom originally occupied by the sulfonate ion. In view of the relative basicities
of HO- and H2O, the hydroxy ligand of the [hydroxy(aquo)iodo]benzene ion (PhI+(OH2)OH) is expected to be
strongly bound and the water ligand is expected to be weakly bound to the iodine(III) center. This species has a
pKA at (4.30 ( 0.05). PhI+(OH2)OH and its conjugate base are present in equilibrium with the [hydroxy(auqo)]-
µ-oxodiphenyldiiodine cation (Ph(HO)I-O-I+(OH2)Ph). This µ-oxo dimer is present at significant levels even in
relatively dilute solutions as the combination equilibrium constant is (540 ( 50). This dimer can be protonated, and
the pKA of the conjugate acid is ≈2.5. The equilibrium constant for dimerization of [oxo(aquo)iodo]benzene
(PhI+(OH2)O-), the most important monomer in acidic solutions, is ≈8.6.
Introduction
heteroligand bond distances in HTIB is not observed with
symmetrical aryl λ3-iodanes such as PhI(OAc)211,12 and PhICl2.13
This difference is attributed to the much greater basicity of the
hydroxide ligand compared with the tosylate ligand.8 From a
structural standpoint, HTIB, and presumably HMIB, bear some
resemblance to iodonium salts, e.g., Ar2I+X-, although they are
substantially more electrophilic at the iodine(III) atom. HMIB
[Hydroxy(mesyloxy)iodo]benzene (1, HMIB) and its tosyloxy
analog (2, HTIB) are stable, crystalline organoiodine(III)
compounds that can be prepared by treatment of (diacetoxy-
iodo)benzene (3) with methanesulfonic acid and H2O in
chloroform1 or acetonitrile,2 or p-TsOH‚H2O in dichloroethane3
or acetonitrile.4 They are electrophilic at iodine and useful for
the phenyliodination and/or oxysulfonylation of a range of
organic substrates.5,6 The structure of HTIB has been estab-
lished by single-crystal X-ray analysis7 and, as expected from
the bonding,8-10 possesses a T-shaped configuration in which
the iodine-oxygen bonds are co-linear (O-I-O ) 178.8°)
and nearly orthogonal (86.0°, 92.8°) to the I-CAr bond. The
I-OH bond (1.94 Å) is shorter than that predicted (1.99 Å) for
covalent radii, while the I-OTs bond (2.473 Å) is much longer
and partially ionic in character. The large difference in iodine-
1
2
3
and HTIB are soluble in water at natural pH and hold promise
as synthetic reagents in this medium. Not only is water an
environmentally benign solvent, but it may also participate in
“solvohyperiodination” reactions.14,15 For example, while ke-
tones are converted to R-mesylate1,16 and R-tosylate17 derivatives
with HMIB and HTIB in organic solvents, reaction 1, they have
been shown to react with HTIB in water to give R-hydroxy
ketones,18 reaction 2. A determination of the actual iodine(III)
* Author to whom correspondence should be addressed.
X Abstract published in AdVance ACS Abstracts, October 1, 1997.
(1) Zefirov, N. S.; Zhdankin, V. V.; Dan’kov, Yu. V.; Koz’min, A. S.;
Chizov, O. S. J. Org. Chem. USSR (Engl. Transl.) 1985, 21, 2252.
(2) Stang, P. J.; Surber, B. W.; Chen, Z.-C.; Roberts, K. A.; Anderson,
A. G. J. Am. Chem. Soc. 1987, 109, 228.
(3) Neiland, O. Ya.; Karele, B. Ya. J. Org. Chem. USSR (Engl. Transl.)
1970, 6, 889.
(4) Koser, G. F.; Wettach, R. H. J. Org. Chem. 1977, 42, 1476.
(5) Moriarty, R. M.; Vaid, R. K.; Koser, G. F. Synlett 1990, 365.
(6) Koser, G. F. [Hydroxy(tosyloxy)iodo]benzene. In Encyclopedia of
Reagents for Organic Synthesis; Paquette, L. A., Ed.; J. Wiley: New York,
1995, Vol. 4, p 2782.
(7) Koser, G. F.; Wettach, R. H.; Troup, J. M.; Frenz, B. A. J. Org.
Chem. 1976, 41, 3609.
(1)
(2)
(8) Koser, G. F. Hypervalent halogen compounds. In The Chemistry of
Functional Groups, Supplement D; Patai, S., Rappoport, Z., Eds.; Wiley:
Chichester, 1983, Chapter 18, p 721; see pp 729-740.
(9) Varvoglis, A. The Organic Chemistry of Polycoordinated Iodine;
VCH Publishers, Inc.: New York, 1992; see pp 4-7.
(10) Koser, G. F.; McConville, D. B.; Rabah, G. A.; Youngs, W. J. J.
Chem. Crystallogr. 1995, 25, 857 and references cited therein.
(11) Lee, C.-K.; Mak, T. C. W.; Li, W.-K.; Kirner, J. F. Acta Crystallogr.
1977, B33, 1620.
species present in aqueous solutions of HMIB and HTIB will
clarify the modes of action of these compounds.
(13) Archer, E. M.; van Schalkwyk, T. G. D. Acta Crystallogr. 1953, 6,
88.
(14) Moriarty, R. M.; Khosrowshahi, J. S.; Prakash, O. Tetrahedron Lett.
1985, 26, 2961.
(15) Moriarty, R. M.; Prakash, O.; Duncan, M. P.; Vaid, R. K.; Musallam,
H. A. J. Org. Chem. 1987, 52, 150.
(16) Lodaya, J. S.; Koser, G. F. J. Org. Chem. 1988, 53, 210.
(12) Alcock, N. W.; Countryman, R. M.; Esperas, S.; Sawer, J. F. J.
Chem. Soc., Dalton Trans. 1979, 854.
S0002-7863(97)01751-4 CCC: $14.00 © 1997 American Chemical Society

Characterization of Species in Aqueous Solutions
J. Am. Chem. Soc., Vol. 119, No. 41, 1997 9615
the stock solution volume changed only 1-2% over the entire
experiment so that in data fits no adjustments were made for solute
concentration changes.
Equilibrium Constants. Equilibrium constant measurements were
made with solutions containing 0.100 mol‚dm-3 of sodium methane-
sulfonate, so that the reported equilibrium constants are concentration
constants for solutions where the ionic strength was 0.1 mol‚dm-3. It
was assumed that activity coefficients of all species participating in
equilibrium were reasonably close to 1.00, since the solutions were
“dilute”.
A number of uncertainties need to be addressed. For
example, do HMIB and HTIB ionize upon dissolution in H2O
to give the hydroxy(phenyl)iodonium ion (PhI+OH) and the
corresponding sulfonate ions (RSO2O-) and, if so, are they
present as ion pairs or fully solvated ions? What are the pKA
values of the >I-OH group of HMIB and HTIB in H2O? Are
they the same, as expected for fully solvated ions, or different?
A variety of µ-oxodiphenyldiiodine(III) compounds of general
structure 4 in which X is typically a nucleofugic group, e.g.,
-OSO2CF3,19,20 -OCOCF3,21 -OClO3,19 or -ONO2,22 are now
known and are formally anhydridesof the putative hydroxy-
Solubilities of HMIB and HTIB in Water. The solubility of HTIB
in water is 0.024 g per mL (61 mmol‚dm-3) at 22 °C.25 The solubility
of HMIB in water was determined to be 0.71 g per mL (2.25 mol‚dm-3
)
at room temperature. The pH of a 2.25 mol‚dm-3 HMIB solution is
about 2. The solubility of HMIB at pH > 4.3 drops significantly. To
obtain solutions with pH greater than this, [HMIB]0 had to be kept at
e3 mmol‚dm-3 to prevent precipitation of iodosylbenzene. Signifi-
cantly, the solubility of HMIB in 1 N NaOH is greater than that
observed near pH 4.3 and under mildly alkaline conditions. The
solubility of HMIB in 1 N NaOH (5.8 mmol‚dm-3) is about twice that
under mildly alkaline conditions.
4
5
6
iodanes, PhI(OH)X. Two µ-oxodiiodine ditosylates, namely,
the heterocyclic species 523 and o-tolyl analog 6,24 are also
known and suggest that equilibrium concentrations of µ-oxo-
diiodine derivatives of HMIB and HTIB may be present in
aqueous solution and contribute to the chemical behavior of
these compounds.
In this report, we use results of UV-vis and NMR spectro-
scopic measurements, along with potentiometric titration data,
to demonstrate the presence of these dimeric µ-oxodiiodine-
(III) species in aqueous solution. We obtain the dimerization
equilibrium constant and the pKA values for the parent mono-
mers. The data presented establish that the [hydroxy(sulfonyl-
oxy)iodo]benzene species 1 and 2 are present in aqueous
solution as fully solvated ions, an organosulfonate ion, and a
free hydroxy(phenyl)iodonium ion. The hydroxy(phenyl)-
iodonium ion produced is the same from either compound and
is essentially hydrated iodosylbenzene in various protonated
forms.
Precipitation of Iodosylbenzene from Solutions of HMIB at pH
> 4.3. We observed that the pH of HMIB solutions could not be
adjusted above ≈4.3 at [HMIB]0 > 3 mmol‚dm-3 or formation of a
cloudy precipitate ensued. A solution of HMIB (0.202 g, 0.64 mmol)
in water (10 mL) was treated with 0.1 N NaOH, sufficient to adjust
the solution to pH 5.31. The solution became cloudy, and after some
time a light yellow precipitate was collected and air dried. The mp
was 201-203 °C, identifying the precipitate as iodosylbenzene, PhIO.
Thermal Decomposition of Aqueous HMIB Solutions. A satu-
rated solution of HMIB (0.7 g, 2.2 mmol) in water (1.0 mL) was
prepared, sealed, and stored in the dark. After several days, white
particles began to appear, and the characteristic aroma of iodobenzene
was detected. After 34 days, the white solid (24.8 mg) was collected
by filtration. The solid was identified as iodylbenzene on the basis of
its melting point (235 °C, explosion) and its solubility in water.26
Aqueous Solutions of Iodylbenzene. PhIO2 was prepared by
oxidation of iodosylbenzene (PhIO) with hypochlorite from Chlorox.
The white solid was recrystallized from water (mp 235 °C, explosion).
The solubility of PhIO2 in water was determined to be about 6.2 mg/
mL (26 mmol‚dm-3). The pH of a saturated solution was 9.68,
corresponding to a pKA of 6.95 for the conjugate acid.
Experimental Materials and Methods
Materials. HMIB and HTIB were synthesized from (diacetoxy-
iodo)benzene as described in the Introduction and purified by recrys-
tallization. All other reagents were purchased reagent grade materials
and were used without further purification.
Spectroscopy and Potentiometric Titrations. UV-vis spectro-
scopic measurements were recorded on a Cary 17 spectrophotometer
updated by OLIS, Inc., Bogart, GA, to allow digital recording of spectra.
All NMR spectra were obtained on a Varian VXR 300 MHz
spectrometer equipped with a broad band probe. Potentiometric
titrations were done by standard methods. Solutions were open to the
air.
Effects of Concentration and pH on UV-Vis Spectroscopy. For
optical spectroscopic measurements, solutions containing the desired
HMIB or HTIB concentration and 0.1 mol‚dm-3 of sodium methane-
sulfonate were prepared. The solution temperature was regulated at
(20.0 ( 0.5) °C, and the pH was adjusted by addition of 1.0 N NaOH.
Quartz cells were filled with solution, and spectrophotometric readings
were taken. Prior to again adjusting the stock solution pH, the solution
from the sample cell was returned to the stock container. In each case,
Results and Discussion
The reactions occuring upon dissolution of HMIB or HTIB
in water are complex. The results are best understood if we
first present the final picture that emerges and then discuss the
results of each set of experiments and how they fit into the
overall picture. The primary processes which occur when
HMIB and HTIB dissolve in water are presented in Scheme 1.
Both λ3-iodanes undergo complete ionization to give the
hydroxy(phenyl)iodonium ion (PhI+OH) and the corresponding
sulfonate ion (RSO2O-) as fully solvated species, i.e., “free”
ions. PhI+OH is presumed to be ligated with at least one water
molecule at an apical site of the iodine(III) atom originally
occupied by the sulfonate ion. In view of the relative basicities
of HO- and H2O, the hydroxy ligand of the [hydroxy(aquo)-
iodo]benzene ion (PhI+(OH2)OH) is expected to be strongly
(25) Wettach, R. H. Ph.D. Dissertation, The University of Akron, May,
1981; see p 63.
(26) Willgerodt, C. Die Organischen Verbindungen mit MehrVertigem
Jod; F. Enke: Stuttgart, 1914.
(27) Smith, R. M.; Martell, A. E.; Motekaitis, R. J. NIST Critical Stability
Constants of Metal Complexes Database, Version 1.0, NIST Standard
Reference Data, Gaithersburg, MD, 1993.
(28) Perdoncin, G.; Scorrano, G. J. Am. Chem. Soc. 1977, 99, 6983.
(29) Moss, R. A.; Alwis, K. W.; Bizzigotti, G. O. J. Am. Chem. Soc.
1983, 105, 681.
(17) Koser, G. F.; Relenyi, A. G.; Kalos, A. N.; Rebrovic, L.; Wettach,
R. H. J. Org. Chem. 1982, 47, 2487.
(18) Luthern, J. M.S. Thesis, The University of Akron, 1986.
(19) Zefirov, N. S.; Zhdankin, V. V.; Dan’kov, Yu. V.; Koz’min, A. S.
J. Org. Chem. USSR (Engl. Transl.) 1984, 20, 401.
(20) Hembre, R. T.; Scott, C. P.; Norton, J. R. J. Org. Chem. 1987, 52,
3650.
(21) Gallos, J.; Varvoglis, A.; Alcock, N. W. J. Chem. Soc., Perkin Trans.
1 1985, 757.
(22) Alcock, N. W.; Countryman, R. M. J. Chem. Soc., Dalton Trans.
1979, 851.
(23) Koser, G. F.; Wettach, R. H. J. Org. Chem. 1980, 45, 1542.
(24) Kalos, A. N. Ph.D. Dissertation, The University of Akron, Jan, 1985;
see pp 143-144.
(30) Moss, R. A.; Boguslawa, W.; Krogh-Jespersen, K.; Blair, J. T.;
Westbrook, J. D. J. Am. Chem. Soc. 1989, 111, 250.
(31) March, J. AdVanced Organic Chemistry, 4th ed.; John Wiley &
Sons: New York, 1992; see pp 250-252.
(32) Schardt, B. C.; Hill, C. L. Inorg. Chem. 1983, 22, 1563-1565.

9616 J. Am. Chem. Soc., Vol. 119, No. 41, 1997
Richter et al.
Scheme 1
H2O
PhI(OH)OSO2R (s)
PhI+(OH2)OH (aq) + RSO3–(aq)
(3)
(4)
(5)
K4
PhI+(OH2)OH
PhI+(OH2)O– + H+
K5
PhI+(OH2)O– + HO–
PhI(OH)O– + H2O
K6
PhI+(OH2)O– + PhI+(OH2)OH
Ph(HO)I–O–I+(OH2)Ph + H+
Ph(HO)I–O–I+(OH2)Ph + H2O (6)
K7
Ph(H2O)I+–O–I+(OH2)Ph
(7)
(8)
K8
Ph(HO)I–O–I+(OH2)Ph
PhIO2 + PhI + H3O+
bound and the water ligand is expected to be weakly bound to
the iodine(III) center.8 Because HMIB and HTIB are largely
insoluble at room temperature in nonhydroxylic solvents such
as CHCl3 and CH3CN, their solubilities in H2O and CH3OH
are attributed to nucleophilic assistance to ionization by solvent
molecules and solvation of the resulting ions.
Figure 2. Dependence of the apparent molar absorbtivity of HMIB
aqueous solutions on pH at (b, O) 300 nm, (2, 4) 320 nm, and (9, 0)
336 nm. Empty symbols are data for 3.0 mmol‚dm-3 HMIB solutions,
while filled symbols are for 6.0 mmol‚dm3 solutions. Lines are
calculated on the basis of the proposed mechanism for 3.0 mmol‚dm-3
solutions, as described in the text.
The dissolution of either HMIB or HTIB in water, reaction
3, Scheme 1, affords 1 molar equivalent of the same weak acid,
namely, the [hydroxy(aquo)iodo]benzene ion (PhI+(OH2)OH),
with pKA ) 4.3. The conjugate base, [oxo(aquo)iodo]benzene
(PhI+(OH2)O-), is a hydrated form of iodosylbenzene [PhsIdO
T Ph-I+-O-) and has limited solubility in water. PhI+(OH2)-
OH and its conjugate base (PhI+(OH2)O-) combine, producing
the [hydroxy(aquo)]-µ-oxodiphenyldiiodine cation (Ph(HO)I-
O-I+(OH2)Ph), reaction 6. This dimer is present at significant
levels even in relatively dilute solutions as K6 ) (540 ( 50).
The initially formed dimer is further protonated in very acidic
pH, giving the [bis(aquo)]-µ-oxodiphenyldiiodine dication (PhI-
(OH2)I+-O-I+(OH2)Ph), which has pKA ≈ 2.5, reaction 7.
Thermal decomposition of solutions of HMIB is slow and
produces iodylbenzene and iodobenzene. The production of
these disproportionation products probably occurs Via a con-
certed decomposition of Ph(HO)I-O-I+(OH2)Ph, reaction 8.
The relative concentrations of iodine(III) species present in
aqueous solutions of HMIB and HTIB are highly pH dependent.
In strongly acidic solutions (epH 1), PhI+(OH2)OH is the
primary monomer present, while Ph(H2O)I+-O-I+(OH2)Ph is
the main dimeric form. In less acidic and moderately basic
Figure 3. Dependence of apparent molar absorbtivity at 320 nm on
pH in aqueous solutions for several [HMIB]0 (mmol‚dm-3): (+, O,
- - -) 0.25; (3, [, - - -) 1.0; (0, 2, - - -) 2.0; (], s) 3.0; (×, ‚-‚) 4.0;
and (+, ‚‚‚, ‚‚‚ -) 5.0. All solutions contained 0.100 mol‚dm-3 of
NaOMs. With the exception of the 5.0 mmol‚dm-3 data, the lines
drawn through the data do not use the equilibrium constants and molar
absorbtivities for the “best fits” (Table 2) but rather are derived from
Scheme 3 using pK4 ) 4.30, K6 ) 540, p(1/K7) ) 2.5, and ꢀih ) 135,
ꢀi ) 140, ꢀdh ) 2860, and ꢀdhh ) 1530 dm3‚mol-1‚cm-1, as discussed
in the text. One of the lines (‚‚‚-) for the 5.0 mmol‚dm-3 data uses
Scheme 2 and the values in Table 2.
solutions (pH 6-11), PhI+(OH2)O- is the dominant monomer
and ultimately separates to give solid iodosylbenzene. At even
higher pH, a new species is generated, presumably PhI(OH)O-,
reaction 5. As expected from the dimerization reaction shown
in reaction 6, [Ph(HO)I-O-I+(OH2)Ph] is maximized in the
pH range near the pKA of PhI+(OH2)OH, i.e., pH 4.3.
In the following sections, results obtained from optical and
NMR spectroscopy, from concentration and pH studies, and
from potentiometric titrations are presented. The supporting
data for inclusion of each reaction in Scheme 1 are presented,
along with arguments for exclusion of others.
Optical Absorption Spectra of Aqueous Solutions of
HMIB: Dependence on pH and Concentration. The optical
absorption spectra of aqueous solutions of HMIB are pH and
concentration dependent, Figures 1, 2, and 3. At all pH, the
significant spectral features occur in the UV, with shoulders
around 275 nm and a broad plateau or slight peak near 330
nm. Only the tails of the UV peaks appear above 400 nm.
Figure 1. Optical absorption spectra of HMIB solutions at various
pH: (‚‚‚) HMIB (2.0 mmol‚dm-3) in 1.0 mmol‚dm-3 methanesulfonic
acid; (s) HMIB (2.0 mmol‚dm-3) in pH 4.25 solution, adjusted by
the addition of NaOH; (- - -) HMIB (2.0 mmol‚dm-3) in pH 9.1,
adjusted by the addition of NaOH; and (‚‚‚) HMIB (4.0 mmol‚dm-3
)
dissolved in 1.0 N NaOH. All solutions, except for the solution with
methanesulfonic acid, had 100 mmol‚dm-3 of sodium methanesulfonate
added. The apparent molar absorbtivity, ꢀ, is the absorbance of the
solution (cm-1), divided by [HMIB]0.

Characterization of Species in Aqueous Solutions
J. Am. Chem. Soc., Vol. 119, No. 41, 1997 9617
Table 1. Molar Absorbtivities (dm3‚mol-1‚cm-1) Used for Fits of
Scheme 2
the Data in Figure 2
H2O
PhI(OH)OSO2R (s)
PhI+(OH2)OH (aq) + RSO3–(aq)
(3)
(4)
nm
ꢀih
ꢀi
ꢀdh
K4
300
320
336
530
135
45
240
140
150
2670
2300
1800
PhI+(OH2)OH
PhI+(OH2)O– + H+
K6
PhI+(OH2)O– + PhI+(OH2)OH
Ph(HO)I–O–I+(OH2)Ph + H2O (6)
Table 2. Values of Equilibrium Constants and Molar
The pH dependence of the apparent molar absorbtivities (ꢀ)
are shown in Figure 2, where ꢀ is the solution absorbance (cm-1
Absorbtivities (dm3‚mol-1‚cm-1) Giving Best Fits of the Apparent
Molar Absorbtivity at 320 nm versus pH Data in Figure 3. All
Measurements Were Taken at (20.0 ( 0.5) °C. All Solutions
Contained 0.10 mol‚dm-3 NaOMs
)
divided by [HMIB]0. Two aspects of the plots at 300, 320,
and 336 nm are striking: maxima appear between pH 4.0 and
4.5, and the pH at which the maximum occurs (pHmax) shifts
with λ. At low pH, the ꢀ appear to approach a constant value.
For 7 < pH < 11, the ꢀ are constant: above pH 11, the ꢀ change,
indicating the appearance of an additional iodine(III) species,
e.g., PhI(OH)O-, reaction 5. The spectrum of the iodine(III)
species at pH 14 is substantially different from that at 7 < pH
< 11, Figure 1. Clearly, the pKA of PhI(OH)O- is very high.
In alkaline solutions, or in strongly acidic solution (pH 1),
the ꢀ are independent of concentration over a significant range,
Figure 3. Thus, the ꢀ are concentration and pH independent in
strongly acid solution and in basic solution, at pH < 11. In
contrast, the ꢀ are strongly concentration dependent in the region
around pH 4.3.
[HMIB]0
pK4
K6
ꢀih
ꢀi
ꢀdh
0.25
0.99
2.0
3.0
4.0
4.29
4.31
4.30
4.29
4.35
4.36
590
620
670
670
690
720
139
124
125
140
137
137
143
132
133
145
145
190
2400
2420
2390
2470
2430
2465
5.0
shifts in pHmax were well predicted, as were the peak profiles.
Significantly, the 6.0 mmol‚dm-3 data deviate from the calcu-
lated curve as pH approaches pK4: this is consistent with
Scheme 2 since the ratio of the concentrations of dimers to
monomers increases as [HMIB]0 increases, while the absorb-
tivities of the µ-oxodiphenyldiiodine species per mole of iodine-
(III) are much larger than those of the monomers. Thus, Scheme
2 provides a good approximation of the equilibria present in
aqueous HMIB solutions.
The wavelength dependence of pHmax, Figure 2, and the
concentration dependence of ꢀ, Figure 3, are consistent with a
pH-dependent dimerization near 4.3. These observations can
be explained with a simple mechanism which includes only
reactions 3, 4, 6, and water dissociation (KW), Scheme 2. Mass-
balance and equilibrium relationships yield eqs 9 and 10:
The system is more complicated than suggested by Scheme
2. For a series of solutions with [HMIB]0 from 0.25 to 5.0
mmol‚dm-3, ꢀ320 was measured as a function of pH, Figure 3.
The data were fit using eqs 9 and 10, much as described above,
except that ꢀih and ꢀi were not constrained to the experimental
values. At each [HMIB]0, after ꢀih and ꢀi were selected, K4
was varied to match pHmax. Finally, K6 and ꢀdh were varied to
match ꢀmax and the PWHH. Table 2 lists the values giving best
fits, which were excellent: to illustrate, the best fit line for the
5 mmol‚dm-3 data is included in Figure 3. ꢀih for the best fits
was essentially constant, and about the same as the experimental
ꢀih. ꢀih most influences the fits at pH < pK4. The best-fit ꢀi
increased somewhat with [HMIB]0 but were close to the
experimental value at low concentrations. ꢀi most influences
ꢀ320 at pH > pK4. The best-fit ꢀdh were nearly constant, and at
low concentrations were close to the best-fit value from Figure
2. The best-fit pK4 increased slightly at [HMIB]0 above 4.0
mmol‚dm-3; however, at lower concentrations, pK4 was nearly
constant and indistinguishable from the value used for Figure
2. K6 showed the most consistent variation with [HMIB]0,
changing from 590 to 720 as the concentration was increased
2K4K6
[H+]
K4
[PhI+(OH2)OH]2 + 1 +
[PhI+(OH2)OH] -
[HMIB]0 ) 0 (9)
+
(
)
[H ]
K4K6 ꢀdh
[H+]
[PhI+(OH2)OH]2 +
ꢀiK4
ꢀih +
[PhI+(OH2)OH] ) A (10)
+
(
)
[H ]
A is the optical absorbance (cm-1) of the sample at a selected
wavelength, and ꢀih, ꢀi, and ꢀdh are the molar absorbtivities of
PhI+(OH2)OH, PhI+(OH2)O-, and Ph(HO)I-O-I+(OH2)Ph,
respectively. Best fits of the data in Figure 2 were obtained
using eqs 9 and 10. Concentrated base was used to adjust the
pH so that the total volume change was only 1-2%, and volume
corrections were unnecessary. The constants whose values must
be assigned are K4, K6, ꢀih, ꢀi, and ꢀdh. At each pH, [PhI+-
(OH2)OH] was calculated from eq 9. A, and hence ꢀ, was
calculated from eq 10.
from 0.25 to 5.0 mmol‚dm-3
. The K6 obtained for 2-3
mmol‚dm-3 solutions, i.e., 670, is the same as that obtained
for 320 nm data for Figure 2, where [HMIB]0 was 3.0
mmol‚dm-1. In summary, the best-fit constants for the data in
Figure 3, based on Scheme 2 and summarized in Table 2, give
the following values: at 320 nm, ꢀih, ꢀi, and ꢀdh are (135 ( 10),
(140 ( 10), and (2400 ( 50) dm3‚mol-1‚cm-1, respectively;
pK4 ) (4.30 ( 0.02); and K6 e 590.
The position of pHmax is controlled by two factors: pK4 and
the relative values of ꢀih and ꢀi. Experimental ꢀih and ꢀi can be
obtained from solutions at pH 1 and pH > 8, respectively,
leaving only pK4 to adjust pHmax. The K6 and ꢀdh selected
control the magnitude of ꢀ at pHmax, i.e., ꢀmax. Increasing either
K6 or ꢀdh increases ꢀmax; however, increasing K6 dramatically
increases the peak width at half-height (PWHH), which is not
true for ꢀdh. Thus, by varying K6 and ꢀdh, both the experimental
ꢀmax and PWHH can be matched. The calculated lines in Figure
2 were obtained with pK4 ) 4.30, K6 ) 670, and [HMIB]0 )
All of the changes in the “constants” for the best fits tabulated
in Table 2 (data in Figure 3) can be attributed to changes in
solution character at higher pH, resulting from increases in
[PhI+(OH2)O-], reaction 4. Iodosylbenzene is only slightly
soluble in water: although supersaturated solutions apperaed
to form and precipitation of the gelatinous product was slow,
we observed that solutions with [HMIB]0 > 3 mmol‚dm-3
became visibly cloudy at pH > 4.7 within several minutes. The
3.0 mmol‚dm-3
.
ꢀdh was varied to match ꢀmax at each
wavelength, and experimental ꢀih and ꢀi were used, Table 1.
All data with [HMIB]0 ) 3.0 mmol‚dm-3 were reasonably fit
using the experimental ꢀih and ꢀi, and the same pK4 and K6.
The ꢀdh giving the best fits varied with λ, as expected. The

9618 J. Am. Chem. Soc., Vol. 119, No. 41, 1997
Richter et al.
Scheme 3
H2O
PhI(OH)OSO2R (s)
PhI+(OH2)OH (aq) + RSO3–(aq)
(3)
(4)
K4
PhI+(OH2)OH
PhI+(OH2)O– + H+
K6
PhI+(OH2)O– + PhI+(OH2)OH
Ph(HO)I–O–I+(OH2)Ph + H+
Ph(HO)I–O–I+(OH2)Ph + H2O (6)
K7
Ph(H2O)I+–O–I+(OH2)Ph
(7)
effect of the light scattering caused by the particulates is
artificially high absorbances, which shift the pHmax to slightly
higher values and broaden the peaks. Thus, artificially large
K6 and ꢀi are needed to fit the data. As a result, the K6 obtained
at the lowest [HMIB]0 is an upper limit, i.e., K6 e 590. The
effect is erratic since the degree of precipitation depends on
the age of the solution, on a time scale of minutes. Thus, in
Figure 2 the molar absorbtivity at pH 11 was the same at 3 and
6 mmol‚dm-3 only because the data were taken with sufficient
rapidity. Additional evidence of light scattering is seen in the
data treatment described next. Please take note: with the
exception of the 5.0 mmol‚dm-3 data, the lines in Figure 3 are
not best fits yielding the values in Table 2 but are derived from
Scheme 3 as described later.
Figure 4. Dependence of ꢀ320 on [HMIB]0 at selected pH. Data points
are taken from the best fits of Figure 2 at six [HMIB]0. Lines are
calculated fits to the data as described in the text: (b) pH 2.98; (×)
pH 3.31; (9) pH 3.64; (+) pH 3.97; (top curve) pH 4.30.
The data in Figure 3 define a three-dimensional surface, where
the axes are pH, [HMIB]0, and ꢀ320. Figure 3 simply shows
slices through this space with [HMIB]0 held constant. Using
the smooth lines through the ꢀ320 vs pH data generated using
the constants in Table 2, we can take slices of pH-[HMIB]0-
ꢀ320 space with pH held constant, giving plots of ꢀ320 vs [HMIB]0
at selected pH, Figure 4. These data cannot be collected easily
in a direct fashion. The motivation for such plots is as follows.
The potentiometric titration data, discussed later, suggest that
protonation of Ph(HO)I-O-I+(OH2)Ph, reaction 7, may be
important in the lower pH range of our measurements. Thus,
the system is better described by Scheme 3, with Kw. Scheme
3 is the same as Scheme 2, with the addition of reaction 7.
Mass-balance and equilibrium relationships give eqs 11 and 12:
Figure 5. Dependence of K (b) and E (2) on [H+]-1 for aqueous
HMIB solutions.
agreement of the values will provide strong support for the
mechanism. The E vs [H+]-1 plot will give (K6ꢀdh) as the slope
and (K6K7ꢀdhh) as the intercept, yielding ꢀdh and ꢀdhh: we have
a value for ꢀdh from other data sets, providing a further test of
the mechanism.
K4
Figure 4 plots were fit as follows. For a selected pH, [PhI+-
(OH2)OH] was calculated from eq 11 for a series of [HMIB]0.
At each [HMIB]0, A was calculated using eq 12, yielding ꢀ320
at that [HMIB]0. Experimental ꢀih and ꢀi were used, Table 1.
The pK4 obtained from the fits of the ꢀ vs pH data at low
[HMIB]0 in Figures 2 and 3 was used (4.30). At each pH, K
and E were varied to obtain the best fit of the ꢀ320 vs [HMIB]0
data for that pH. K determines the curvature of the initial
portion of the curve, while E determines the slope at higher
concentrations. Excellent fits were obtained with this model,
Figure 4. As seen in Figure 5, the K vs [H+]-1 and E vs [H+]-1
plots were straight lines for data from pH 3.0 to 4.6, a change
in [H+] by a factor of almost 50. The linearity of these two
plots lends considerable support to the equilibria in Scheme 3.
A linear regression fit of the K vs [H+]-1 plot gives K6 )
(434 ( 7) and (1/K7) ) (1.2 ( 0.3) × 10-3. Significantly, this
value of K6 is not very different from that obtained from Figure
2 (K6 ) 670) using Scheme 2, or the K6 from Figure 3 (K6 e
590) using Scheme 2. Using the K6 and K7 obtained here, a
linear regression fit of the E vs [H+]-1 data give ꢀdh and ꢀdhh
2K4K[PhI+(OH2)OH]2 + 1 +
[PhI+(OH2)OH] -
[HMIB]0 ) 0 (11)
+
(
)
[H ]
K4
K4E[PhI+(OH2)OH]2 + ꢀih + ꢀi
[PhI+(OH2)OH] ) A
+
(
)
[H ]
(12)
(13)
where
1
K t K6 K7 +
+
(
)
[H ]
and
ꢀdh
E t K6 ꢀdhhK7 +
(14)
+
(
)
[H ]
ꢀdhh is the molar absorbtivity of the [bis(aquo)]-µ-oxodiphenyl-
diiodine dication (Ph(H2O)I+-O-I+(OH2)Ph) formed in reac-
tion 7. Now, we have good values for pK4, ꢀi, and ꢀih: at each
pH, we could vary K and E to obtain the best fits of the ꢀ320 vs
[HMIB]0 plots. Equations 13 and 14 now provide an excellent
test of Scheme 3: if the mechanism is correct, a plot of K vs
[H+]-1 will give a straight line, and a plot of E vs [H+]-1 would
also be linear. The K vs [H+]-1 plot would yield K6 as the
slope and (K6K7) as the intercept, allowing calculation of both
K6 and K7: we already have a limiting value for K6 and
equal to (3000 ( 70) and (1500 ( 700) dm3‚mol-1‚cm-1
,
respectively. The ꢀdh obtained is slightly higher than that
obtained from fits of Figures 2 and 3 using Scheme 2, i.e., 2400
and 2300 dm3‚mol-1‚cm-1, respectively. The deviations in K6
and ꢀdh are not so large when we consider the large difference
in the way the values were obtained.
It is significant to note that when K and E at pH > 4.6 are
included, the plot of K vs [H+]-1 curves downward, while the

Characterization of Species in Aqueous Solutions
J. Am. Chem. Soc., Vol. 119, No. 41, 1997 9619
plot of E vs [H+]-1 curves upward. Both trends result from
precipitation of iodosylbenzene and the consequent light scat-
tering. This curvature probably also affects the equilibrium
constants and molar absorbtivities obtained. For example, when
K vs [H+]-1 was fit for pH 3.0-4.0, the K6 obtained was (487
( 14), closer to the value of e590 obtained from Figure 3. We
can reasonably conclude that the actual K6 lies somewhere
between these, i.e., K6 ) 540 ( 50. For p(1/K7), 2.5 was found
vs 2.9 for 3.0 e pH e 4.6. In this smaller pH range, E vs
[H+]-1 gave ꢀdh ) 2860 dm3‚mol-1‚cm-1, closer to those
obtained from fits of Figures 2 and 3. ꢀdhh was 1530
dm3‚mol-1‚cm-1
.
To evaluate the consistency of the data fits, theoretical lines
based on the above model were calculated for Figure 3: the
best fits of the data sets in this figure were used to generate the
three-dimensional surface yielding Figure 4. The theoretical
lines were calculated using a single set of equilibrium constants
and molar absorbtivities: pK4 ) 4.30, K6 ) 540, p(1/K7) )
2.5, and ꢀih ) 135, ꢀi ) 140, ꢀdh ) 2860, and ꢀdhh ) 1530
dm3‚mol-1‚cm-1. Good agreement is seen between the theo-
retical lines and the data, Figure 3, except at higher pH and
higher [HMIB]0, where precipitation of iodosylbenzene produces
deviations.
Figure 6. Optical absorption spectrum of Ph(HO)I-O-I+(OH2)Ph,
calculated as described in the the text (‚-‚): the values plotted are ꢀdh
÷ 3. Spectra of the pH 1 (‚‚‚) and pH 9 (- - -) monomers, from Figure
1, are included for comparison.
Spectra of HMIB and HTIB solutions in the visible region
showed only the trailing edge of the UV absorptions. HMIB
in 1 mol‚dm-3 methanesulfonic acid produced no visible color,
even at concentrations as high as 120 mmol‚dm-3. At natural
pH, a 64 mmol‚dm-3 HMIB solution was light yellow, with
the corresponding increased molar absorbtivity relative to the
acid solution in the range from 450 to 550 nm. When NaOH
was added to the natural pH HMIB solution in a 1:2 molar ratio
(NaOH:HMIB), the yellow color of the solution visibly intensi-
fied, as reflected by strongly increased molar absorbtivities.
Moderately alkaline (pH 8-9) or strongly alkaline (1 N NaOH)
solutions were colorless. In virtually every respect, equivalent
solutions of HMIB or HTIB gave identical results.
Spectrum of the Dimeric [Hydroxy(aquo)-µ-oxodiphenyl-
diiodine Cation. The optical absorption spectrum of Ph(HO)I-
O-I+(OH2)Ph was calculated from spectra of HMIB solutions
at pH 1, pH 9.6, and pH 4.25. The mechanism predicts that
for a 2 mmol‚dm-3 solution at pH 1 or 9.6 the spectrum is due
almost entirely to PhI+(OH2)OH or PhI+(OH2)O-, respectively.
The spectrum at pH 4.25 given in Figure 1 was obtained from
a 2.00 mmol‚dm-3 solution. Using eqs 11 and 13, based on
Scheme 3, the concentrations of PhI+(OH2)OH, PhI+(OH2)O-,
Ph(HO)I-O-I+(OH2)Ph, and Ph(H2O)I+-O-I+(OH2)Ph can
be calculated for this pH and [HMIB]0. Using pK4 ) 4.30, K6
) 540, and p(1/K7) ) 2.5, the respective concentrations of the
iodonium species are 0.762, 0.679, 0.280, and 0.005 mmol‚dm-3.
About 28.5% of the total iodine(III) species are present in
dimeric form. Then, to a good approximation, the molar
absorbtivity at a given wavelength of Ph(HO)I-O-I+(OH2)Ph
is
Figure 7. Potentiometric titrations of HMIB aqueous solutions: (s)
calculated curve based on Scheme 3 and values for the K’s given in
the text; (‚‚‚) calculated curve based on 1 mol of strong acid per mole
of HMIB; (- - -) calculated curve based on 1 mol of weak acid per
mole of HMIB with pKA ) 4.3. 3 mmol‚dm-3 HMIB with (O) 0
mmol‚dm-3 NaOMs, (0) 100 mmol‚dm-3 NaOMs, (4), or 200
mmol‚dm-3 NaOMs. (×) HTIB, 3 mmol‚dm-3 in water with no other
additives.
the value of 2860 dm3‚mol-1cm-1 obtained from the series of
data fits described previously.
Potentiometric Titrations of Aqueous HMIB Solutions.
Potentiometric titrations of HMIB solutions showed the presence
of 1 equiv of titratable protons per mole of HMIB. The shape
of the titrations curves does not conform to that of a strong
acid, nor to that of a simple monoprotic weak acid. The addition
of dimer formation to the solution dynamics, reaction 6, accounts
for the large slopes seen in the titration curves, Figure 7. The
difference between the potentiometric titration curves of strong
acids, weak acids, and weak acids with a dimerization equilib-
rium in the region of the pKA is dramatic, so that it is clear
which type of species is present. The large slope in the pKA
region of the titration curve can be considered a “fingerprint”
for the presence of PhI+(OH2)OH and related species. Initial
fits using a model including only reactions 3, 4, 6, and KW,
i.e., Scheme 2, showed that the pKA was near 4.3. Using a
model including reactions 3, 4, 6, 7, and KW, i.e., Scheme 3,
the potentiometric titration data were very well fit, Figure 7.
The data fits were substantially improved by inclusion of the
protonation equilibrium for the initially formed dimer, reaction
7. In the fits of the titration data, the value chosen for K4, the
KA of PhI+(OH2)OH, essentially varies the y-axis offset and
hence the y-axis intercept of the curve. The value chosen for
ꢀ - 0.381ꢀih - 0.340ꢀi
ꢀdh ) 2
(15)
(
)
0.285
The calculated spectrum is given in Figure 6, along with the
spectra of PhI+(OH2)OH and PhI+(OH2)O-. The molar ab-
sorbtivities of the dimer in the region from 300 to 400 nm are
dramatically larger than those of the monomers, consistent with
assertions that the yellow color of aqueous solutions of a number
of iodine(III) species results from the -I-O-I- unit of the
µ-oxodiphenyldiiodine dimer.33 The calculated value of ꢀdh at
320 nm is 2485 dm3‚mol-1cm-1, in reasonable agreement with
(33) Dasent, W. E.; Waddington, T. C. Proc. Chem. Soc. 1960, 71.

9620 J. Am. Chem. Soc., Vol. 119, No. 41, 1997
Richter et al.
Table 3. 1H Chemical Shifts of the Phenyl and Methyl Protons of
HMIB, HTIB, and NaOMs in D2O at Natural pH
K6, the dimerization equilibrium constant, increases the slope
by dropping the initial points in the titration and raising the
later points, with a “pivot point” near 0.5 mequiv of base per
mmol of HMIB. The solid line in Figure 7 was calculated using
pK4 ) 4.30, K6 ) 540, and p(1/K7) ) 2.5, values obtained from
the treatments of the dependence of optical molar absorbtiVity
on pH. The essential agreement of the data with the calculated
line, based on equilibrium constants obtained from optical
measurements, substantiates the proposed mechanism and
solution species.
ortho-1H
para-1H
meta-1H
methyl-1H
HMIB
HTIB
NaOMs
8.192
8.163
7.751
7.725
7.593
7.566
2.785
2.357
2.795
slightly more than 1 equiv of neutralizable protons was obtained.
Although the number of equivalents of neutralizable protons
per mole of HMIB in the potentiometric titrations was very
nearly equal to one, in many titrations this number was
significantly greater than one: values as high as 1.07 were
obtained. Clearly, the production of “extra” acid is related to
the thermal decomposition of HMIB since prompt titrations after
solution preparation gave 1.00 ( 0.01 equiv per mole. Since
PhIO2 itself is a weak base (the conjugate acid has a pKA of
6.95), and since we expect that PhI would be neutral in its
reaction with water, it would seem that the “extra acid” in aged
solutions arises from an interaction of products and reactants
not heretofore recognized.
NMR Spectroscopy of HMIB and HTIB Solutions. Po-
tentiometric titrations of aqueous HMIB and HTIB solutions
indicate that identical aqueous species are produced from these
iodine(III) compounds. 1H NMR spectra in D2O support this
conclusion. At natural pH, the peaks due to the ortho-, meta-,
and para-hydrogens of the iodine(III)-bound phenyl group of
HMIB and HTIB occur at the same positions, Table 3.
Furthermore, the methyl protons of the methanesulfonate species
in HMIB solutions give chemical shifts identical to those seen
in D2O solutions of sodium methanesulfonate, a result expected
only if the methanesulfonate ions in the HMIB solutions are
free from the iodine(III) species.
For illustrative purposes, the potentiometric titration curves
expected for production of 1 mol of strong acid per mole of
HMIB, and for 1 mol of weak acid with pKA ) 4.30 per mole
of HMIB, are included in Figure 7. These visually demonstrate
how different the behavior of HMIB solutions are from those
of simple strong or weak acids. The clear deviation of the
potentiometric titration curves from the simple cases eliminates
some possible reactions following dissolution of HMIB. For
example, it is not possible that the PhI+(OH2)OH formed in
reaction 3 immediately and completely deprotonates, giving
PhI+(OH2)O-. If this were so, then addition of HMIB to water
would produce 1 equiv of strong acid (H+) and 1 mol of a weak
acid with a pKA of 4.3, PhI+(OH2)O-, per mole of HMIB, and
this was not observed. We can conclude that pKA observed at
4.3 is not that of PhI+(OH2)O-. Thus, the titration results
support a mechanism in which the species exhibiting a pKA at
4.3 is the initially formed [hydroxy(aquo)iodo]benzene cation.
The three sets of data plotted in Figure 7 are for different
concentrations of added sodium methanesulfonate (0, 0.100, and
0.200 mol‚dm-3). The coincidence of the plots shows a virtual
lack of dependence of the system on ionic strength, and lack of
dependence on [MsO-], in the range examined. The lack of
dependence on [MsO-] implies that the equilibrium constant
for formation of ion pairs between MsO- and PhI+(OH2)OH is
either very large (essentially complete ion-pairing at a one-to-
one concentration ratio) or very small (not detectable with
[MsO-] ) 0.200 mol‚dm-3). The absence of a large equilibrium
constant for ion-pairing is seen from the following experiment.
When an aqueous solution of 3.00 mM HTIB in water was
titrated in the same manner as the HMIB solutions at 20 °C,
virtually identical results were obtained as with HMIB solutions,
Figure 7. This is consistent with a model in which HMIB and
HTIB both dissolve in water to give organosulfonate ions and
identical aqueous iodine(III) species.
The presence of the various equilibria shown in Scheme 1
are manifest in the NMR spectra. Solutions of iodobenzene in
CD3OD show well-developed fine structure in the peaks for its
aromatic protons, while the corresponding 1H peaks for HMIB
in D2O or CD3OD appear virtually as two simple broadened
triplets and a broadened doublet (iodobenzene is insoluble in
water). The proton of the hydroxy ligand initially present in
HMIB and HTIB clearly is exchanged to the HDO pool, i.e.,
there is no proton signal corresponding to an -OH group other
than HDO.
Monomeric Species. Optical spectroscopy and potentio-
metric titrations of aqueous solutions of HMIB or of HTIB
establish that between pH 1 and 10 there are two monomeric
iodine(III) species present, i.e., the [hydroxy(aquo)iodo]benzene
ion and its conjugate base. Identical solution species are
produced from the two compounds. Optical absorbance results
for solutions above pH 11 show that a new species appears there.
The initially formed monomer is the [hydroxy(aquo)iodo]-
benzene cation (PhI+(OH2)OH) formed when the organosul-
fonato ligand is displaced by a water molecule. In this species,
the H2O ligand is loosely bound to the iodine(III) center,8 while
the hydroxide ligand is covalently bound. The lack of signifi-
cant covalent character in the water-iodine(III) bond is rational-
ized as follows. If the bond were purely covalent, the positive
charge would reside on the oxygen atom, and the species would
be a strong acid. Witness the pKA values of other tricovalent
oxygen-centered compounds: the pKA of H3O+ is -1.74,27
while the pKA of protonated dimethyl ether, (CH3)2O+-H is
Thermal Stability of HMIB Solutions. The observed
products of the thermal decomposition of HMIB in H2O are
the disproportionation products iodylbenzene and iodobenzene.
Disproportionation of HMIB most likely occurs via heterolysis
of a µ-oxodiphenyldiiodine species such as Ph(HO)I-O-I+-
(OH2)Ph, where decomposition is initiated by deprotonation of
the hydroxy group, reaction 8:
(8)
The only significant divergence between the calculated and
measured data in the potentiometric titrations occurred beyond
the point where 1.00 equiv of base per mole of HMIB had been
added. It was regularly observed that if the titrations were done
within 30-60 min after solution preparation, then 1.00 ( 0.01
equiv of acid was obtained per equivalent of HO-; however, if
some time elapsed from solution preparation to titration, then
-2.5.28 Thus, a species in which the water ligand is covalently
¨
bound to the iodine(III) would be very acidic, with a pKA many
units below 4.3. This would result in the production of 1 mol
of strong acid and 1 mol of weak acid (pKA ) 4.3) per mole of
HMIB dissolved, and this was not the case. Thus, we conclude
that the water ligand(s) is (are) weakly bound to the iodine-

Characterization of Species in Aqueous Solutions
J. Am. Chem. Soc., Vol. 119, No. 41, 1997 9621
Table 4. pKA Values for Heterocyclic Hydroxy(aryl)iodanes with
Internal Carboxylate and Alkoxide Ligands
This lowers the degree of positive charge at iodine and decreases
the acidity of the hydroxyl function. The greater acidity of the
bis(trifluoromethyl)benziodoxole 9 compared to 8 and the
greater acidity of 10 compared with 11 is consistent with this
logic. Even if the iodanes 7 and 8 ionize in H2O and are ligated
with one or more water molecules they would not give free
ions, since the anions would be held in close proximity to the
iodonium centers, reaction 17. Both HMIB and HTIB give the
(17)
free [hydroxy(aquo)iodo]benzene cation, PhI+(OH2)OH, upon
dissolution in aqueous solution. The high acidity of this species
(pKA ) 4.3) is understandable if it is the hydroxy ligand and
not the water ligand, which deprotonates. The water ligand is
a much weaker base than either the carboxylate and alkoxide
ligands in 7 and 8, respectively, so that we would expect a pKA
substantially below 7, such as 4.3.
What is necessary to obtain [hydroxy(B)iodo]arenes in which
the B:- ligand is covalently bound to the iodine(III)? The
discussion in the preceding paragraph provides insight. The
observed trend is that an increase in the basicity of B:- lowers
the acidity of the proton in >I-O-H; this is because the >I-B
bond is less ionic in character. The least basic of the B:- ligands
in Table 4 is in compound 7, which has the most acidic -OH
group. The conjugate acid of a typical carboxylate ion has a
pKA of about 4.8. Bases whose conjugate acids have pKA values
greater than this should be able to form covalent bonds in
[hydroxy(B)iodo]arenes. The conjugate acid of H2O is H3O+,
which has a pKA of -1.74.31 H2O is a very weak base relative
to the bases known to form covalent bonds: it should be no
surprise that the bond with H2O in PhI+(OH2)OH is of a very
low order. Clearly, there is an upper limit to the basicity of
B:- which can be tolerated by the [3c-4e] bond, which must be
polar with significant electron density residing on the axial
ligands. For example, [hydroxy(methoxy)iodo]benzene has not
been isolated; the pKA of CH3OH is 15.2.31 [Bis(methoxy)-
iodo]benzene has been isolated, but it decomposes explosively.32
Thus, we see that a pKA of 4.3 is not unreasonable for
deprotonation of the hydroxy ligand. Now consider the relative
stabilities of the possible conjugate bases of PhI+(OH2)OH, i.e.,
PhI+(OH2)O- and PhI(OH)2. Just as [hydroxy(methoxy)iodo]-
benzene cannot be isolated due to instability resulting from the
very high basicity of the two axial ligands, we would expect
PhI(OH)2 to be of a low stability. In contrast, PhI+(OH2)O- is
a hydrate of PhIO, iodosylbenzene, a very stable molecule. Thus,
it is highly unlikely that PhI(OH)2 is the conjugate base of PhI+-
(OH2)OH. The structural considerations discussed all support
the assignment of PhI+(OH2)O- as the conjugate base, while
the water ligand remains very loosely bound. The probable
identity of the high pH monomeric species is PhI(OH)O-:
(III), with a bond order significantly less than one. Single-
crystal X-ray analysis of HTIB7 shows that the HO-I bond length
is 1.94 Å, consistent with a covalent bond between the iodine-
(III) atom and oxygen, while the TsO-I bond length is 2.47 Å,
indicating significant ionic character in this iodine-oxygen
bond. Thus, both solution and solid state data suggest that the
structure forms a stabilized unit with any additional ligands
exhibiting low bond order.
The hydroxide ligand in PhI+(OH2)OH is a very weak base,
i.e., the conjugate acid formed in reaction 16, PhI+(OH2)2, is a
strong acid. Were this not the case, the natural pH of HMIB/
HTIB solutions would be alkaline, whereas they are acidic.In
PhI+(OH2)OH + H2O a PhI2+(OH2)2 + HO- (16)
the pH range of our experiments, we could not discern the effects
of the putative PhI2+(OH2)2 dication. In summary, it is PhI+-
(OH2)OH which has an acid dissociation constant with pKA )
4.3. Given that H2O is loosely bound, and that HO- is tightly
bound, what can we say about the deprotonation that results in
the observed pKA at 4.3?
pKA measurements of acyclic hydroxy(aryl)-λ3-iodanes have
not been reported. In contrast, heterocyclic hydroxy(aryl)-
iodanes with internal carboxylate and alkoxide ligands are
known, and pKA values for a few of them have been re-
ported,29,30 Table 4. It is noteworthy that the acidity of the >I-
O-H moiety in these hydroxyiodanes is inversely related to
the basicity of the second heteroligand. For example, the
benziodoxolone 7 is at least 5600 times more acidic than the
dimethylbenziodoxole 8 and manifests the lower basicity of an
internal carboxylate ligand compared to that of an internal
alkoxide ligand. Presumably, ionic structures of ArIOHB
contribute less to the hybrid as the basicity of B:- increases,
i.e.,
PhI+(OH2)O- + HO- a PhI(OH)O- + H2O
(5)
In the above discussion, we examined the results and
theoretical arguments, reaching the conclusion that the three
monomeric forms of aqueous iodosylbenzene are PhI+(OH2)-
OH, PhI+(OH2)O-, and PhI(OH)O-. HMIB is extremely water

9622 J. Am. Chem. Soc., Vol. 119, No. 41, 1997
Richter et al.
soluble (0.71 g/mL, 2.25 mol‚dm-3) because the MsO- ligand
structures suggest that they are formed via interactions of the
anions with dimeric species already present in solution. In
moderately acid solution, Ph(HO)I-O-I+(OH2)Ph is the main
dimer: anions such as nitrate can displace the water ligand and
17 precipitates:
is highly water soluble. HTIB is less soluble (0.024 g/mL, 0.061
mol‚dm-3 25 due to the lesser solubility of TsO- (relative to
)
MsO-) resulting from the aromatic ring in this anion. Thus,
HMIB solutions (and the MsO- counterion) provide greater
insight into the solubilities of the three monomers. If the pH
is <2.3, then >99% of the monomer present is PhI+(OH2)OH.
Thus, the “solubility of HMIB at natural pH” gives a lower
limit on the solubility of PhI+(OH2)OH: it is Very soluble,
>2.25 mol‚dm-3. This solubility is consistent with an ionic
species that is easily solvated. At pH >5.3 through mildly
alkaline conditions, >90% of the monomer is PhI+(OH2)O-:
this monomer is soluble only to the extent of about 3
mmol‚dm-3. Such low solubility is consistent with the zwit-
terionic character of this species, and with the observation that
dehydration produces PhIO, a highly stable uncharged molecule.
In highly alkaline pH, PhI+(OH2)O- deprotonates to PhI(OH)O-.
This species is more soluble than PhI+(OH2)O- and dissolved
at about 6 mmol‚dm-3 at the highest pH tested.
Ph(HO)I-O-I+(OH2)Ph + NO3- f
Ph(HO)I-O-I(NO3)Ph V + H2O (19)
In very acidic solution, the bis(aquo) dimer, Ph(H2O)I+-O-
I+(OH2)Ph, predominates: nitrate can replace both water
ligands, and 16 is obtained:
Ph(H2O)I+-O-I+(OH2)Ph + 2NO3- f
Ph(NO3)I-O-I(NO3)Ph V + 2H2O (20)
17
15
16
Dimerization in Acidic Media. The equilibrium constant
for dimerization of PhI+(OH2)OH to produce the [bis(aquo)]-
µ-oxodiphenyldiiodine dication (Ph(H2O)I+-O-I+(OH2)Ph),
reaction 18,can be calculated since this reaction is the sum of
Zhdankin and co-workers34,35 prepared the tetrafluoroborate,
hexafluoroantimonate, and hexafluorophosphate disalts of oxy-
bis[phenyliodonium], 18a-c, by treating chloroform solutions
of (diacetoxyiodo)benzene with aqueous solutions of the cor-
responding acids at room temperature. Oxybis[phenyl-, bis-
(tetrafluoroborate(1-))] iodonium (18a) was obtained (70%)
from reaction of (diacetoxyiodo)benzene with aqueous HBF4
and was the first such salt reported.34 Hydroxy-µ-oxodiphe-
nyldiiodine(1+), tetrafluoroborate(1-) (19) was obtained when
18a was treated with water for 2-3 h. In summary, synthetic
results support our conclusions regarding the presence of dimeric
iodine(III) species in solution and the presence of more than
one protonated form.
2PhI+(OH2)OH a Ph(H2O)I+-O-I+(OH2)Ph + H2O (18)
reactions 4, 6, and 7. Thus, K18 ) K4K6K7 and K18 ≈ 8.6. HMIB
is very soluble in water at natural pH (2.25 mmol/mL). At pH
≈1, PhI+(OH2)OH is the primary monomer: at 0.1 mol‚dm-3
,
about 50% of the iodine(III) species will be present as Ph(H2O)-
I+-O-I+(OH2)Ph. Such concentrations of the [hydroxy-
(aquo)]-µ-oxodiphenyldiiodine cation cannot be achieved since
this dimer forms primarily near pH 4.3 (i.e., pK4) where
[PhI+(OH2)O-] become significant, and the solubility of
PhI+(OH2)O- is only ≈3 mmol‚dm-3
.
Acyclic Hydroxy(aryl)-λ3-iodanes, Their Anhydrides, and
Aqueous Solutions of [Hydroxy(sulfonyloxy)iodo]arenes.
Apart from the [hydroxy(sulfonyloxy)iodo]arenes 12, acyclic
hydroxy(aryl)-λ3-iodanes are rare. For example, stable acyclic
hydroxyiodanes such as 13 and 14, containing carboxylate or
nitrate ligands, have not been isolated. However, dimeric
µ-oxodiphenyldiiodine species such as [bis(acetato-O)]-µ-oxo-
diphenyldiiodine (15), formally anhydrides of these putative
hydroxyiodanes, have been isolated.
19
18
Iodosylbenzene in the solid state is an amorphous powder,
where in the PhIO units form polymeric chains.36 Probably,
each chain contains one water of hydration so that the polymer
has HO endcaps, i.e.,
Ph
HO
I
O
H
n
From this point of view, [bis(aquo)]-µ-oxodiphenyldiiodine is
the diprotonated dimer of iodosylbenzene, while [hydroxy-
(aquo)]-µ-oxodiphenyldiiodine is the monoprotonated dimer of
iodosylbenzene. One or both of the HO endcaps can be replaced
by various anions, such as nitrate, producing [bis(nitrato-O)]-
µ-oxodiphenyldiiodine or [hydroxy(nitrato-O)]-µ-oxodiphenyl-
diiodine, respectively. From this same point of view, we can
consider HMIB to be an adduct of methanesulfonic acid with
the monomeric PhIO. We can consider
12
13
14
Dasent and Waddington33 obtained 15 by treatment of
(diacetoxyiodo)benzene with nitric acid. (Diacetoxyiodo)-
benzene is substantially soluble in nitric acid, while crystals of
15 slowly precipitate from solution. Willgerodt,26 as later shown
by Dasent and Waddington,33 obtained [bis(nitrato-O)]-µ-
oxodiphenyldiiodine (16), upon treatment of iodosylbenzene
with nitric acid. On the other hand, when aqueous solutions
of HTIB were treated with sodium nitrate solutions, [hydroxy-
(nitrato-O)]-µ-oxodiphenyldiiodine (17) was obtained by Wet-
tach.25 The absence of hydroxyiodanes, PhI(OH)X, correspond-
ing to these dimers and the pH dependence of the dimeric
O–
Ph I+
a special unit of exceptional stability.
Conclusions
(34) Zhdankin, V. V.; Tykwinski, R.; Caple, R.; Berglund, B.; Koz’min,
A. S.; Zefirov, S. S. Tetrahedron Lett. 1988, 29, 3717.
(35) Zhdankin, V. V.; Tykwinski, R.; Berglund, B.; Mullikin, M.; Caple,
R.; Zefirov, N. S.; Koz’min, A. S. J. Org. Chem. 1989, 54, 2609.
(36) Carmalt, C. J.; Crossley, J. G.; Knight, J. G.; Lightfoot, P.; Martin,
A.; Muldowney, M. P.; Norman, N. C.; Orpen, A. G. J. Chem. Soc., Chem.
Commun. 1994, 2367.
The primary processes which occur when HMIB and HTIB
dissolve in water are presented in Scheme 1. Both λ3-iodanes
undergo complete ionization to give the hydroxy(phenyl)iodo-
nium ion (PhI+OH) and the corresponding sulfonate ion
(RSO2O-) as fully solvated species, i.e., “free” ions. PhI+OH

Characterization of Species in Aqueous Solutions
J. Am. Chem. Soc., Vol. 119, No. 41, 1997 9623
is presumed to be ligated with at least one water molecule at
an apical site of the iodine(III) originally occupied by the
sulfonate ion, and this water molecule is loosely bound.
Solution behavior is strongly influenced by the presence of
µ-oxodiphenyldiiodine dimeric species. The monomeric solu-
tion species from HMIB and HTIB are simply various proto-
nated states of hydrates of iodosylbenzene. Many of the
reactions described herein will also be important in solutions
in organic solvents in which water has not been rigorously
excluded.
Laboratory of the University of Notre Dame. Substantial
intellectual contributions were made by Prof. K. C. Calvo of
the Chemistry Department at the University of Akron. Samples
of HMIB and HTIB were prepared by Roger A. Moore. We
also acknowledge the essential support of this project by the
Molecular Spectroscopy Laboratory of the University of Akron,
and in particular to Takishi Saito for his instruction of one of
us (T.D.Z.) in the operation of the instrumentation. Financial
support was provided by the University of Akron. Chemical
structures and reactions were prepared by Ms. Charlotte A.
Bogart using ChemDraw or ISIS/Draw.
Acknowledgment. This work was begun during the sab-
batical leave of the corresponding author at the Radiation
JA971751C