Reactions of molybdenum(vi) with metal ion reductants

Article abstract of DOI:10.1039/b516547b

The reactions of aqueous H₂MoO₄ at low pH with titanium(ii), titanium(iii), europium(ii), vanadium(ii), and germanium(ii), as monitored at 430 nm, give biphasic profiles featuring a sharp rise in absorbance followed by a marked decrease (Fig. 1). The final product is the dimeric Mo(v) cation, [Mo₂O₄]²⁺, and the strongly absorbing intermediate is taken as a monomeric Mo(v) species. The molar absorbances of the transients from different reductants are not the same, nor are the rate laws governing the fadings. None of the decay curves exhibits evidence of a second order dependence on the transient. The kinetic behaviors of these systems are consistent with the intervention of successor complexes of the type, (formed by inner sphere reductions of Mo(vi)), which decompose, via first-order processes, to a monomeric Mo(v) species. The latter then experiences rapid dimerization, which is kinetically silent. The possibility that Ge(ii) bypasses the unstable tripositive state by reducing Mo(vi) to Mo(iv) (which then undergoes rapid Mo(vi)-Mo(iv) comproportionation) is considered. The Royal Society of Chemistry 2006.

Full text of DOI:10.1039/b516547b

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www.rsc.org/dalton | Dalton Transactions
Reactions of molybdenum(VI) with metal ion reductants†‡
Zhiyong Yang and Edwin S. Gould*
Received 21st November 2005, Accepted 10th March 2006
First published as an Advance Article on the web 19th April 2006
DOI: 10.1039/b516547b
The reactions of aqueous H2MoO4 at low pH with titanium(II), titanium(III), europium(II),
vanadium(II), and germanium(II), as monitored at 430 nm, give biphasic profiles featuring a sharp rise
in absorbance followed by a marked decrease (Fig. 1). The final product is the dimeric Mo(V) cation,
[Mo2O4]2+, and the strongly absorbing intermediate is taken as a monomeric Mo(V) species. The molar
absorbances of the transients from different reductants are not the same, nor are the rate laws governing
the fadings. None of the decay curves exhibits evidence of a second order dependence on the transient.
The kinetic behaviors of these systems are consistent with the intervention of successor complexes of
the type
, (formed by inner sphere reductions of Mo(VI)), which decompose, via first-order
processes, to a monomeric Mo(V) species. The latter then experiences rapid dimerization, which is
kinetically silent. The possibility that Ge(II) bypasses the unstable tripositive state by reducing Mo(VI)
to Mo(IV) (which then undergoes rapid Mo(VI)–Mo(IV) comproportionation) is considered.
Stoichiometric studies
Introduction
The biphasic nature of the reactions studied made stoichiometric
Although the diverse role of molybdenum derivatives in redox
catalysis is extensively recognized, past interest has been centered
mainly around oxygen atom transfers from peroxy and oxo
compounds to alkenes, sulfides, and amines, as promoted by
molybdate centers. Such reactions are generally considered to
involvecomplexes ofMo(VI) and the oxidant.2 Catalysis of electron
transfer between metal centers in aqueous media is less usual.
It has recently been noted that reductions of 1,4-quinones,
vanadium(IV) and triiodide, by both di- and tripositive titanium
are catalyzed by Mo(VI), and evidence has been offered implicating
the pentapositive state, Mo(V), in such reactions. However, the
dimeric cation, [Mo2O4]2+, a persistent form of Mo(V) in the range
0.01–2.0 M H+,3 has been shown to be devoid of catalytic activity
in such systems.
determinations unfeasible at or near 430 nm, the region of
principal interest. However, information concerning the overall
conversions could be obtained at 710 nm in 11.0 M HCl, where
Mo(V) has been shown to be converted to the stable green complex,
[MoOCl5]2−.8 Measured, deficient quantities of Na2MoO4 were
added to a known excess of the reductant. Increases in absorbance
were compared with those from addition of excess Mo(VI). The
results are summarized in Table 1.
Kinetic studies
Reactions were performed under argon, using freshly prepared
solutions of the reductants. Rates were evaluated from changes
in absorbance at 430 nm, using a Durrum–Gibson stopped flow
spectrophotometer in conjunction with an OLIS computer system.
Temperatures were kept at 22.0 0.5 ◦C. The ionic strength was
maintained with NaCl/LiCl/HCl or CF3SO3Na/CF3SO3H. All
runs were carried out with the reducing agent in four- to tenfold
excess.
The present study deals with the generation of the catalytically
active Mo(V) species and its conversion to the inactive dimeric
form.
Experimental
Materials
In each case, a marked increase in absorbance, followed by
a sharp decrease, was observed (for example, Fig. 1). Peak
absorbance generally occurred within the first second after mixing,
but with Eu(II) the very rapid initial rise was barely discernable.
With the other reactants in excess, profiles were analyzed by
treating the transformations as consecutive pseudo-first order
processes (1),
All solutions were prepared using Millipore-Q system deionized
water that had been boiled for 2 h and then purged with argon for
a further 2 h to remove dissolved oxygen. Titanium metal wire,
titanium(III) chloride, sodium molybdate, and trifluoromethane-
sulfonic acid (triflic acid) (all Aldrich reagents) were used as
received. Preparations of V(II),4 and Ti(II)5 (used as triflate salts)
and Eu(II)6 and Ge(II)7 (used as chlorides) were as described
elsewhere.
k
k
2
1
Mo(VI) + Red −→ intermediate −→ products
(1)
then, using the procedure of Bose9 to disentangle the contributing
rate constants (k1 and k2), and to estimate the extinction coefficient
of the intermediate. For the Mo(VI)–Eu(II) reaction, k1 could not
be obtained.
Department of Chemistry, Kent State University, Kent, Ohio 44242, USA
† Electron transfer. Part 1631.
‡ Electronic supplementary information (ESI) available. Tables S-1 to S-6:
detailed kinetic data for redox reactions.
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Table 1 Generation of molybdenum(V) from H2MoO4. Stoichiometric studiesa
Reductant
Ti(II)
[Red.]/mM
[H2MoO4]/mM
[Mo(V)]/mM
[Mo(V)]/[Red.]
2.5
5.0
5.3
6.0
4.0
6.6
3.0
4.5
10.0
15.0
10.0
10.0
9.0
11.0
11.0
11.0
5.0
10.0
5.3
6.2
4.1
6.6
6.2
9.2
2.0
2.0
1.00
1.03
1.02
1.00
2.1
Ti(III)
Eu(II)
V(II)
Ge(II)
2.0
a Reactions carried out in 11 M HCl: k = 710 nm. (The product is [MoOCl5]2−.)
Generation of the Mo(V) intermediate
Kinetic curves (e.g., Fig. 1) are analyzed as a superposition
of two consecutive pseudo-first order processes.10 One of these
components arises from the growth of an intermediate species,
the other from its disappearance. The usual ambiguity associated
with the assignment of rate constants from such components10
does not apply here, since one of the rate constants in each system
is proportional to the concentration of the reductant, whereas the
other, in all cases except Ti(III), is independent of the reductant
concentration.
Two of the rate laws for formation of the intermediate are
straightforward monomial expressions:
rate = k [Mo(VI)][Red.] kV(II)(2.3 0.1) × 102M−1s−1
(7)
k
Ge(II)(9.4 0.8) × 102M−1s−1
Fig. 1 Kinetic profile for the reaction of H2MoO4 (0.30 mM) with V(II)
(12 mM), [H+] = 0.25 M; l = 0.50 M, (CF3SO3H–CF3SO3Na); k = 430 nm;
T = 22 ◦C. The solid line represents the experimental curve; the circles were
calculated by non-linear least squares refinement utilizing the relationship
of Bose (ref. 9a in text), kinetic parameters in Tables 2 and 3, and einterm in
Table 4.
In contrast, reductions by Ti(II) (Table S-1‡) and Ti(III)
(Table S-2‡) are appreciably inhibited by H+, and that by Ti(III) is
also retarded slightly by chloride. Data for intermediate formation
by Ti(II) conform to eqn (8) and that by Ti(III) to eqn (9).
[Mo(VI)][Ti(II)]k KA
(rate)Mo,Ti(II)
=
(8)
KA + [H+]
Results
[Mo(VI)][Ti(III)]k KA[H+]
In the spectral region of interest (ca. 430 nm), neither H2MoO4 nor
the predominant Mo(V) product, Mo2O42+(aq) (kmax = 295 nm),
absorbs appreciably, but measurements at 710 nm in 11 M
HCl (Table 1), where Mo(V) is converted to the green complex
(rate)Mo,Ti(III)
=
(9)
(1 + KA[Ti(III)](1 + KA[H+] + KCl[Cl−])
Refinement of data for the formation of the intermediate by
Ti(II) and Ti(III) yields the kinetic parameters listed in Table 2.
Calculated and observed rate constants (in parentheses) are
compared in Tables S-1 and S-2‡.
2−
8
MoOCl5
, confirm the quantitative formation of pentavalent
molybdenum.
The generation of the intermediate formed by the action of
Eu(II) was too rapid to measure.
2 Mo(VI) + Ti(II)) → 2 Mo(V) + Ti(IV)
(2)
(3)
(4)
(5)
(6)
Decay of the Mo(V) intermediate
Mo(VI) + Ti(III) → Mo(V) + Ti(IV)
Mo(VI) + Eu(II) → Mo(V) + Eu(III)
2 Mo(VI) + V(II) → 2 Mo(V) + V(IV)
2 Mo(VI) + Ge(II) → 2 Mo(V) + Ge(IV)
With one exception, rate laws for the decomposition of the
intermediate do not include concentrations of the redox partners.
The transient formed in the Mo(V)–Ge(II) reaction is consumed
exponentially (k = 0.40 0.03) s−1. Decays of those formed from
the Mo(VI)–Ti(II) and the Mo(VI)–V(II) systems (Tables S3 and
S4‡) are retarded at higher acidities, with both decompositions
proceeding in accordance with rate law (eqn (10)),
kKA[Int]
KA + [H+]
−d[Int]/dt =
(10)
Note, however, that under the conditions used in the kinetic runs
(reductant in excess), both Ti(II) and V(II) would be expected to
yield the respective tripositive states.
with the k value for V(II) almost three times as great as that
pertaining to Ti(II) (Table 3).
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Table 2 Generation of Mo(V) intermediate. Rate laws and kinetic parametersa
Red.
Medium
Rate laws
Parameters
V(II)
Ge(II)
CF3SO3H/CF3SO3Na
HCl/LiCl
k[Mo(VI)][V(II)]
k[Mo(VI)][Ge(II)]
[Mo(VI)][Ti(II)]kKA
KA + [H+]
[Mo(VI)][Ti(III)]kKA[H+]
(1 + KA[Ti(III)])(1 + KA[H+] + KCl[Cl−])
k = (2.3 0.1) × 102 M−1s−1
k = (9.4 0.8) × 102 M−1s−1
Ti(II)
CF3SO3H/CF3SO3Na
k = (6.2 0.3) × 103 M−1s−1
KA = (0.23 0.02) M
Ti(III)
HCl/LiCl
k = (150 20) M−1s−1
KA = (0.44 0.02) M
KCl = (0.7 0.3) M−1
k > 1 × 104 M−1s−1b
Eu(II)
HCl/LiCl
a Reactions at 22.0 0.5 ◦C, l = 0.50 M, k = 430 nm. b Lower limit, assuming rate = k [Mo(VI)][Eu(II)].
Table 3 Decomposition of Mo(V) intermediate. Rate laws and kinetic parametersa
Red.
V(II)
Medium
Rate lawsb
Parameters
kKA[Int]
KA + [H+]
CF3SO3H/CF3SO3Na
k = (5.7 0.9) s−1
KA = (0.15 0.04) M
k = (2.1 0.05) s−1
kKA[Int]
KA + [H+]
Ti(II)
CF3SO3H/CF3SO3Na
KA = (0.42 0.02) M
Ge(II)
Ti(III)
HCl/LiCl
HCl/LiCl
k[Int]
k = (0.41 0.03) s−1
k[Ti(III)][H+][Cl−][Int]
k = (8.7 0.7) × 103 M−3 s−1
KTi = (2.1 0.3) × 102 M−1
k = (21 1) s−1
1 + KTi[Ti(III)]
(k + kꢀ[Cl−])[Int]
Eu(II)
HCl/LiCl
(1 + [H+]/KHA + KCl[Cl−])
kꢀ = (58 4) M−1 s−1
KCl = (0.29 0.15) M−1
KHA = (0.22 0.04) M
a Reactions at 22.0 0.5 ◦C; l = 0.50 M, k = 430 nm. b Rates refer to decomposition of the intermediate, Int, −d[Int]/dt.
Table 4 Extinction coefficients of Mo(V) intermediate (430 nm)a
The intermediate from the Eu(II) reduction is consumed accord-
ing to a two-component rate law (eqn (11)),
Red.
103 eInt/M−1 cm−1
(k + kꢀ[Cl−])[Int]
−d[Int]/dt =
(11)
(1 + [H+]/KHA + KCl[Cl−])
Ti(II)
Ti(III)
V(II)
Ge(II)
Eu(II)
1.12 0.11
1.54 0.28
1.50 0.23
0.84 0.02
0.33 0.02
in which the denominator reflects partition of the intermediate
into three species; protonated, deprotonated (inactive) and a Cl−
addition complex (active).
Among the systems examined, only the Mo(VI)–Ti(III) reaction
involves a transient that decays with the participation of the parent
reductant. The rate law (eqn (12))
a e Values obtained by disentangling the biphasic reaction curves using the
method of Bose (ref. 9a).
k[Ti(III)][H+][Cl−][Int]
−d[Int]/dt =
(12)
1 + KTi[Ti(III)]
Discussion
exhibits kinetic saturation, suggesting formation of a Ti(III)–
intermediate adduct (Kformation = 210 M−1).
The array of biphasic reaction profiles observed in this study
indicates that Mo(VI) reacts quickly in acid solutions with each
of our metal-center reductants, yielding an intermediate of lower
oxidation state which strongly absorbs near 430 nm and which
decays ultimately to the familiar cation, [Mo2O4]2+. The properties
of the intermediate correspond to the metastable monomeric
Mo(V) species proposed by Sykes11 and generated electrolytically
by Chalipoyil and Anson.12 This monomer is converted to a more
stable chloro complex in 10 M HCl8 but decomposes quickly
in 2 M CF3SO3H,13 and appears to be involved also in the
Note that none of the decay curves display evidence of a kinetic
component that is second order in the disappearing intermediate.
Table 4 lists the molar absorptivity of the transient species,
calculated from the disentangling of each of the reaction profiles.9a
Extinction coefficients for intermediates in the reductions by Ti(III)
and by V(II) are in fair agreement, but those obtained in the
remaining systems are quite different, with the mismatch well
outside the estimated uncertainty. The implication is that we are
dealing with four different absorbing species.
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Mo(VI)-catalyzed reductions of benzoquinone, vanadium(IV), and
tri-iodide by hypovalent titanium.1
The various rate laws for decomposition of the Mo(V) transients
(Table 3) provide information concerning the speciation of the suc-
cessor complexes. With both the V(II) → V(III) and Ti(II) → Ti(III)
systems, dissociation (17) requires deprotonation, (as indicated by
[H+] in the denominators), whereas the Eu(III) intermediate exists
in both protonated (reactive) and nonprotonated (inactive) forms,
and a modestly stable chloro-complex (KCl = 0.3 M−1) as well.
The Ti(III) → Ti(IV) intermediate, alone among the successors
listed, requires action of additional reductant for dissociation,
as well as both protonation and Cl− attachment. The kinetic
saturation with respect to Ti(III) noted here allows us to estimate a
stability constant near 200 M−1 for coordination of the additional
reductant. The possibility of interaction between Ti(III) and
Ti(IV) centers calls to mind the formation of a Ti(III)–Ti(IV)
complex in aqueous sulfate solution as reported by Fraser and
coworkers.18
The generation of Mo(V) by both V(II) and Ge(II) proceeds
according to straightforward monomial rate laws (Table 2), but
both conversions pose mechanistic questions. The Mo(VI)–Ge(II)
reaction may avoid the unstable state Ge(III) by preliminary
formation of Mo(IV), probably via oxo transfer (13), followed by
rapid Mo((IV),(VI)) comproportionation (14).
(13)
(14)
The rate constant for the Mo(VI)–V(II) reaction, 2.30 ×
102 M−1s−1, exceeds the 1 × 102 M−1s−1 upper limit imposed by
slow substitution at the V(II) center,4,14 thus indicating either a
predominant outer-sphere path or an inner-sphere path involving
substitution at the Mo(V) center for this transfer.
Acknowledgements
The Mo(VI)–Ti(II) rate law (eqn (8)) mirrors partition of
the reductant into two protonation levels, of which only the
deprotonated form is active, and brings to mind the reduction of
chlorinated quinine, chloranilic acid,15 by Ti(II). In contrast, the
denominator of the rate law for the Mo(VI)–Ti(III) reaction, (eqn
(9)), indicates partition of Ti(III) into three species: deprotonated
and protonated forms and a chloro-complex. Of the three, only
[Ti(OH)]2+ is significantly active, a feature common of many
reductions by this center.16
The second phase of our profiles, a sharp absorbance decrease
in all cases (Table 3), is taken here to reflect the transformations of
the monomeric Mo(V) transient to the inactive dimer, [Mo2O4]2+.
Two features of these curves are quite unexpected: (A) The decay
portions exhibit no sign of second order character, which one
would expect for a monomer-to-dimer conversion; and (B) The
calculated molar absorbances of transients generated from the
several reductants do not match (Table 4).
We therefore conclude that the losses of the transients are first
order processes preceding a rapid, kinetically silent, dimerization,
and that these transients, although related, are not the same
species. Moreover, the rate laws and kinetic parameters associated
with the decays are not the same (Table 3).
The kinetic behavior of these intermediates is consistent with
the intervention, in at least some cases, of binuclear successor
complexes that are formed in the initial electron transfer (15) and
which subsequently dissociate into a monomeric Mo(V) species
(17) before rapid conversion to the dimeric product (18):17
We are grateful to the National Science Foundation for partial
support of this work and to Mrs Arla Dee McPherson for technical
assistance.
References
1 Part 162: Z. Yang and E. S. Gould, Dalton Trans., 2005, 1781.
2 See, for example: A. Keller, J. M. Sobczak and J. J. Ziolkowski, in
Molybdenum; An Outline of its Chemistry and Uses, ed. E. R. Braith-
waite and J. Haber, Elsevier, Amsterdam, 1994, ch. 11.
3 (a) Y. Sasaki, R. S. Taylor and A. G. Sykes, J. Chem. Soc., Dalton
Trans., 1975, 396; (b) D. E. Linn, Jr., S. K. Ghosh and E. S. Gould,
Inorg. Chem., 1989, 28, 3225; (c) M. Ardon and A. Pernick, Inorg.
Chem., 1973, 12, 1484.
4 J. C. Chen and E. S. Gould, J. Am. Chem. Soc., 1973, 95, 5198.
5 U. Kolle and P. Kolle, Angew, Chem., Int. Ed., 2003, 42, 4540.
6 F.-R. Fan and E. S. Gould, Inorg. Chem., 1974, 13, 2639.
7 O. A. Babich and E. S. Gould, Inorg. Chem., 2000, 39, 4119.
8 (a) See, for example:H. Diebler and Y. C. Millan, Bol. Soc. Chil. Quim.,
1984, 29(2), 281; (b) A. A. Bergh and G. P. Haight, Jr., Inorg. Chem.,
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4239.
9 (a) R. N. Bose and E. S. Gould, Inorg. Chem., 1985, 24, 2832; (b) R. G.
Wilkins, The Study of Kinetics and Mechanism of Reactions of Transition
Metal Complexes, Allyn and Bacon, Boston, MA, 1974, pp. 20–24.
10 See, for example, J. H. Espenson, Chemical Kinetics and Reaction
Mechanisms, 2nd Edn, McGraw-Hill, New York, 1995, ch. 4.
11 G. R. Gayley, R. S. Taylor, R. K. Wharton and A. G. Sykes, Inorg.
Chem., 1977, 16, 1377.
12 P. Chalipoyl and F. C. Anson, Inorg. Chem., 1978, 17, 2418.
13 M. T. Paffett and F. C. Anson, Inorg. Chem., 1981, 20, 3967.
14 B. R. Baker, M. Orhanovic and N. Sutin, J. Am. Chem. Soc., 1967, 89,
722.
(15)
(16)
(17)
15 Z. Yang and E. S. Gould, Dalton Trans., 2005, 1781.
16 (a) See, for example:M. Orhanovic and J. E. Earley, Inorg. Chem., 1975,
14, 1478; (b) A. H. Martin and E. S. Gould, Inorg. Chem., 1976, 15,
1934.
17 In view of the recognized slow substitution at the V(II) center,14 step
(15) for the V(II) reduction may be taken to proceed via substitution at
the more labile Mo(VI) center
Mo(VI) + V(II)–OH2 → Mo(VI)–OH–V(II) + H+.
(19)
18 R. T. M. Fraser, R. G. Miller and V. W. Cope, Coord. Chem. Rev., 1966,
1, 85.
(18)
3430 | Dalton Trans., 2006, 3427–3430
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