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The preponderance of inner-sphere paths for peroxodiphosphate ion oxidation of metal complexes
INORG.
NUCL.
CHEM.
LETTERS
Voh
7,
pp.
1-8,
197l.
Pergamon P r e s s .
The Preponderance of Inner-sphere Paths for Peroxodiphosphate
Printed in Great Britain.
Ion Oxidation
of Metal Complexes Eleanor Chaffee
Inge I. Creaser
John O. Edwards
Metca]f Chemical Laboratory Brown University Providence, Rhode Island 02912 (Received lO August 1970)
In our study of oxidations by peroxodiphosphate found for an outer-sphere mechanism. of peroxodisulfate.
ions, no evidence was
This is in contrast to the reactions
The purpose of this eo~munication is to show how very
different are the behaviors of these isoelectrouic peroxides towards metal complexes.
Exemplary data are discussed and a tentative explanation is
given. It was shown (I) that the reaction between Fe(phen)32+ and P2084- (2) in neutraIj aqueous solutions ~#~7~'~'s~_~.,the rate law
_d[FeL~ 2+] : klk2[Fe L32+] [P2054-] dt
k I[L] + kg[F205~-]
suggesting the mechanism Pe L 3
Fe
where L represents,
k! 2 + -~
L 2 2+ +
~ k
FeL 2
2+
+ L
-i
P2084- ~2
Products
in this instance~ a phenanthroline ligamd.
At high con-
centrations of P2084- (above 2 x IO-3M) and in the absence of added phena~:throline, the rate b~=eo_.~s ..... independent of [ P2O8 4- ] and the rate law reduces to
d[FeL32+] = kl[FeL32+] dt
.
No reaction for 14 days at 60 ° in the presence of EDTA.
I mol-lsec -1, pH 1.1-3.5.
This work
c.
d.
c
b.
5 x lO -2 (25 ° )
29-1
No reaction for 14 days at 60 ° .
2+
6.1 x lO -5 (35 °, pH 5.9)
30.6
30.9
Ea(kcal mol "I)
a.
VO
2+
2+
Os(phen) 3
Os (b zpy) 3
2+
1.0 x 10-4 (25 °)
Fe(bipy)32+
Fe(terpy) 2
3.8 x 10 -4 (35 ° )
k (sec-I)
Fe(phen)32+
Complex
_
+z3.4
+24
+26
AS~(cal mol'ideg -I)
Rates and Activation Parameters for the Oxidation bY-P208
TABLE I
12
d
Ref
#
o
o
o
m
Vol. 7, No. 1
PEROXODIPHOSPHATE ION OXIDATION
3
The rate constant for this reaction together with the activation parameters now found are given in Table I.
For comparison the corresponding data for
the acid dissociation are given in Table II.
Since loss of !igand is pos-
tulated to be the rate-determining step in both processes, the parameters should be and are in agreement. The oxidation of Fe(bipy)32+ by P2084- was found to follow the same rate law. Except
at pH extremes~ the rates of dissociation of Fe(phen)32+
and Fe(bipy)32+ are only slightly influenced by pH (4~5).
As expected
from the K_I step the oxidation rates are slower when free ligand is added to the reaction mixtures.
Rates and activation parameters are given in
Tables I and II. T.~J3LE I! Rates and Activation Parameters for the Dissociation
k (sac-1)
Complex
Ea(kcal mol "I)
AS~(deal_~l-i
Fe(phen)32+
3.7 x 10-4 (34.6 ° )
32.1
+28
Fe(bipy)3 2+
1.32 x 10-4(25 °)
28.4
+17
Fe(terpy)22+
I.i x 10 -6 (35 °, pH6)
28.7
7.3 x 10-3 (35 ° , 3 M H +)
28.7
Os(bipy) 3
2+
Os(phen)32+
extremely slow at i00 °
extremely slow at I00 °
Dissociation of the complex Fe(terpy)22+ having terdentate ligands is slower than for the previous two complexes.
In the Fe(terpy]22+ case both
Ref.
5,6
4
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. 1
_
dissociation and oxidation by P208 and If).
are dependent on pH.
(See Tables I
The rate of displacement of a complete ligand was observed at pH 6
at an exchange study while a limiting rate corresponding to the breaking of one Fe-N bond of the tridentate ligand was observed in 3 M H + (7).
The
oxidation rate was found as predicted to lie between the two limits. That an oxidation does take place in the iron complexes rather than a mere
displacement of a ligand for P208
tests:
4-
has been demonstrated by several
SCN- tests for Fe 3+ were positive and the Fe(phen)32+ color
returned after addition of $204
2-
The presence of PO 4
2-
radicals was
supported by the facts that isopropyl alcohol was shown to depress the rate of oxidation of Fe(terpy)2
2+
, Fe(phen)3
2+
and Fe(bipy)3
2+
by P208
4-
Also the induction period was shorter and the rate of polymerization of aerylonitrile faster in the presence of Fe(phen)32+ on Fe(bipy)32+ and
P2084- than in pure P2084-. These results are in marked contrast to those from the reactions between $208
2-
and the same complexes.
The rate law for the initial reac-
tions between FeLx 2+ (L = phen, bipy, terpy) and S2082- has been found by several authors (8,9,10) to be of the form
"d[FeLx] - k [ F e L ] dt
[$2082-]
with rate constants which usually are faster than replacement and that are generally independent at the nature of the ligand.
Some data are sho~m in
Table III. The very inert 0s(bipy)3 2+ did not react with P2084- at 60 ° for at least 14 days. S208
2-
In contrast to this the reaction between Os(bipy)32+ and
is a fast second-order reaction (Table III).
When 0 s'~phen)3 2+ was treated with took place (k ~lO -7 sec "l) initially.
P2084- at 60 ° very slow reaction After a few days the rate increased
and a limiting rate of k = %10 -6 sec -1 (independent of [P2082-]) was reached When a trace of EDTA was added no reaction was observed for at least 14 days
Vol. 7, No. 1
P E R O X O D I P H O S P H A T E ION OXIDATION
5
TABLE III Rates and Activation Parameters for the Oxidation by $208
k(l mol-lsec -I)
Complex Fe(phen)32+
Fe(bipy)3
2+
Fe(terpy)22+
0s(bipy)3
2+
Ea(keal mol -I)
2-
~S~(cal mol-ldeg -I) Ref
0.108 (25 ° )
13.5
-19
8
0.41 (25 °)
10.3
-28
8
0.85 (35 ° )
10.7
-24
8
49.0 (25 ° )
VO 2+
9
a
ii
+
a.
Reaction only measurable in the presence of Ag .
after which time the reaction started and accelerated with time.
This
behavior suggests that a small concentration of an Os(lll) product is _oi:~=u which is able to catalyze the r~action in the absence of EDTA. Ag + and Cu 2+ catalyzed the reaction strongly.
Even [Cu2+] of i0 -~ M had
While the reaction between Os(phen)32+ and S2082- has
a measurable effect.
not been examined it is expected to be similar to the reaction between
Os(bipy)32+
and $208
2-
~ith a similar rate constant.
Prel~minary reactions of P208
4-
with [Os terpy bipy CI] + suggest that
the rate of hydrolysis of the monodentate CI- ion an~ the rate of oxidation with [P2084-] are similar.
The oxidation reaction is strongly catalyzed by
trace metal ions and the rates are only reproducible in the presence of EDTA. 2A ha!f-life of tl/2 50 hr was observed at pH 7, 60°- Oxidation ~y S208 is extremely fast even at room temperature. Attempts at studying the oxidation of [Ru(phen)2PyCl] + and
6
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. 1
[Ru(phen)2pyC~CN] 2+ by P2084- led only to indications of behavior.
The
reaction of the chloro complex was complicated by hydrolysis of reactant and product.
However, it was established that the complex reacts with
P208 ~- (tl/2 % 1.5 hr at pH 9, 60°) at a faster rate than hydrolysis of #
the con~lex (tl/2 ~ 40 P~ at pH 9, 60o) •
For [Ru(phen)2pyCKjCN] 2+,
hydrolysis is too slo~v to interfere with the rate of oxidation.
However,
the solutions turned out to be photosensitive and half-lives of oxidation varied from
tl/2=
0.5 - 18 hr.
The complications found for Ru(II)oxida-
i]ions are too severe to justify firm mechanistic conclusions. While the reaction of vanadyl ion and peroxodisulfate requires the presence of a silver ion catalyst to proceed at room temperature peroxodiphosphate
(ll) the
reacts at a convenient rate with the vanady! ion.
The
kinetics are flrst-order each in oxidant and reductant with a rate constant of 5 x 10 -2 ! mo! -I sec -I at 25 ° , at pH I.i to 3.5. (12) These results, the different kinetics, the great difference in rate constants, and the very different activation parameters,
clearly demonstrate
that the isoelectronic $2082- and P2084- react with these complexes by quite different mechanisms.
While S208
2-
apparcnt!yprefers
an outer-sphere
mechanism the results from the P2084- reactions (except with the R~(II) complexes) are best explained by an inner-sphere mechanism where" the ratedetermining step is the partial dissociation of a ligand with P2084- or 2PO 4 approaching the vacant position follo~Ted by a fast oxidation reaction. In the Os(phen)32+ and Os(bipy)32+ cases no measurable dissociation takes place and so no oxidation is observed as long as catalysts are absent. Tne oxidation potential for the peroxodisulfate co,~le is -2.01 v. (13) The oxidation potential for the peroxodiphosphate has not been determined but the fact that P2084- oxidizes Ag +, Cu 2+ ~nd V02+ indi.cates that the potential is not ~ich different from that of the $208~C coup] e. It is of interest to esti~ate the minimum activation ener~q[ necessary
PEROXO~PHOSPHATE IoN OXIDATION
VoL 7, No. 1
to explain the marked difference in rates of oxidation of Os(phen)32+ by the two isoelectronic peroxides. difference,
Roughly correcting for the temperature
it is found that $2082- oxidizes the osmium complex at least
108 more rapidly than does P2084-. ii kcal mole -1 in E a.
This is equivalent to a difference of
This seems too larae to attribute to a bond disso-
ciation energy in the peroxide bond. explanation.
There is one attractive possible
If the peroxodiphosphate
is considerably more solvated, thel~
desolvation prior to insertion between the ligand planes (as appears to be the case with peroxodisulfate as oxidant) would be required.
The energy
difference could be rationalized as due to about three hydrogen bonds.
Acknowledg/ments.
We thank the USAF Office of Scientific Research for
continuing support of our research.
i.
Sr. A. A. Green, J. O. Edwards and P. Jones, Inorg. Chem., ~, 1858 (1966).
2.
For the purposes of clarity, peroxodiphosphate will be written as P208
irrespective of protonation state.
This does not alter the basic conclusions.
The state of protonation as a function of pH is reported elsewhere. 3.
M.M.
(3)
Crutchfield and J. O. Edwards, J. Amer. Chem. Soc., 82, 3533
(196o). 4.
F. Basolo, J. C. Hayes and H. M. Neumarm, J. Amer. Chem. Soc., 763
3807 (1954). 5.
j. I~irgess and R. H. Prince, J. Chem. Sot., 6C61 (1965).
6.
R. Kogg and R. O. Wilkin.;, J. Ch~m. Soc., 341 (1962).
7.
R. Farina, R. Hogg and R. G. Wilkins, Inorg. Chem., 7, 170 (1968).
8.
J. Burgess and R. H. Prince,
9-
D. H. Irvine, J. Chem. Soc., 2977 (1959).
iO.
J. Chem. Scc., (A), 1772 (1966).
S. Raman and C. H. Brubaker, Jr., J. Inorg. Nucl. Chem., 31, i091 (1969)
8
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. I
Ii.
D. M. Yost and W. H. Claussen, J..~mer. Chem. Soc., 53, 3349 (1931).
13.
W. H. Latimer, "Oxidation Potentials", Prentice-Hall, Inc., Englewood
Cliffs, New Jersey, 2nd Edition (1952), page 78. 12.
M. Andersen, J. 0. Edwards, Sr. A. A. Green, and Sr. M. D. Wis~ell,
Inorg. Chim. Acta, 3, 655 (1969).
NUCL.
CHEM.
LETTERS
Voh
7,
pp.
1-8,
197l.
Pergamon P r e s s .
The Preponderance of Inner-sphere Paths for Peroxodiphosphate
Printed in Great Britain.
Ion Oxidation
of Metal Complexes Eleanor Chaffee
Inge I. Creaser
John O. Edwards
Metca]f Chemical Laboratory Brown University Providence, Rhode Island 02912 (Received lO August 1970)
In our study of oxidations by peroxodiphosphate found for an outer-sphere mechanism. of peroxodisulfate.
ions, no evidence was
This is in contrast to the reactions
The purpose of this eo~munication is to show how very
different are the behaviors of these isoelectrouic peroxides towards metal complexes.
Exemplary data are discussed and a tentative explanation is
given. It was shown (I) that the reaction between Fe(phen)32+ and P2084- (2) in neutraIj aqueous solutions ~#~7~'~'s~_~.,the rate law
_d[FeL~ 2+] : klk2[Fe L32+] [P2054-] dt
k I[L] + kg[F205~-]
suggesting the mechanism Pe L 3
Fe
where L represents,
k! 2 + -~
L 2 2+ +
~ k
FeL 2
2+
+ L
-i
P2084- ~2
Products
in this instance~ a phenanthroline ligamd.
At high con-
centrations of P2084- (above 2 x IO-3M) and in the absence of added phena~:throline, the rate b~=eo_.~s ..... independent of [ P2O8 4- ] and the rate law reduces to
d[FeL32+] = kl[FeL32+] dt
.
No reaction for 14 days at 60 ° in the presence of EDTA.
I mol-lsec -1, pH 1.1-3.5.
This work
c.
d.
c
b.
5 x lO -2 (25 ° )
29-1
No reaction for 14 days at 60 ° .
2+
6.1 x lO -5 (35 °, pH 5.9)
30.6
30.9
Ea(kcal mol "I)
a.
VO
2+
2+
Os(phen) 3
Os (b zpy) 3
2+
1.0 x 10-4 (25 °)
Fe(bipy)32+
Fe(terpy) 2
3.8 x 10 -4 (35 ° )
k (sec-I)
Fe(phen)32+
Complex
_
+z3.4
+24
+26
AS~(cal mol'ideg -I)
Rates and Activation Parameters for the Oxidation bY-P208
TABLE I
12
d
Ref
#
o
o
o
m
Vol. 7, No. 1
PEROXODIPHOSPHATE ION OXIDATION
3
The rate constant for this reaction together with the activation parameters now found are given in Table I.
For comparison the corresponding data for
the acid dissociation are given in Table II.
Since loss of !igand is pos-
tulated to be the rate-determining step in both processes, the parameters should be and are in agreement. The oxidation of Fe(bipy)32+ by P2084- was found to follow the same rate law. Except
at pH extremes~ the rates of dissociation of Fe(phen)32+
and Fe(bipy)32+ are only slightly influenced by pH (4~5).
As expected
from the K_I step the oxidation rates are slower when free ligand is added to the reaction mixtures.
Rates and activation parameters are given in
Tables I and II. T.~J3LE I! Rates and Activation Parameters for the Dissociation
k (sac-1)
Complex
Ea(kcal mol "I)
AS~(deal_~l-i
Fe(phen)32+
3.7 x 10-4 (34.6 ° )
32.1
+28
Fe(bipy)3 2+
1.32 x 10-4(25 °)
28.4
+17
Fe(terpy)22+
I.i x 10 -6 (35 °, pH6)
28.7
7.3 x 10-3 (35 ° , 3 M H +)
28.7
Os(bipy) 3
2+
Os(phen)32+
extremely slow at i00 °
extremely slow at I00 °
Dissociation of the complex Fe(terpy)22+ having terdentate ligands is slower than for the previous two complexes.
In the Fe(terpy]22+ case both
Ref.
5,6
4
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. 1
_
dissociation and oxidation by P208 and If).
are dependent on pH.
(See Tables I
The rate of displacement of a complete ligand was observed at pH 6
at an exchange study while a limiting rate corresponding to the breaking of one Fe-N bond of the tridentate ligand was observed in 3 M H + (7).
The
oxidation rate was found as predicted to lie between the two limits. That an oxidation does take place in the iron complexes rather than a mere
displacement of a ligand for P208
tests:
4-
has been demonstrated by several
SCN- tests for Fe 3+ were positive and the Fe(phen)32+ color
returned after addition of $204
2-
The presence of PO 4
2-
radicals was
supported by the facts that isopropyl alcohol was shown to depress the rate of oxidation of Fe(terpy)2
2+
, Fe(phen)3
2+
and Fe(bipy)3
2+
by P208
4-
Also the induction period was shorter and the rate of polymerization of aerylonitrile faster in the presence of Fe(phen)32+ on Fe(bipy)32+ and
P2084- than in pure P2084-. These results are in marked contrast to those from the reactions between $208
2-
and the same complexes.
The rate law for the initial reac-
tions between FeLx 2+ (L = phen, bipy, terpy) and S2082- has been found by several authors (8,9,10) to be of the form
"d[FeLx] - k [ F e L ] dt
[$2082-]
with rate constants which usually are faster than replacement and that are generally independent at the nature of the ligand.
Some data are sho~m in
Table III. The very inert 0s(bipy)3 2+ did not react with P2084- at 60 ° for at least 14 days. S208
2-
In contrast to this the reaction between Os(bipy)32+ and
is a fast second-order reaction (Table III).
When 0 s'~phen)3 2+ was treated with took place (k ~lO -7 sec "l) initially.
P2084- at 60 ° very slow reaction After a few days the rate increased
and a limiting rate of k = %10 -6 sec -1 (independent of [P2082-]) was reached When a trace of EDTA was added no reaction was observed for at least 14 days
Vol. 7, No. 1
P E R O X O D I P H O S P H A T E ION OXIDATION
5
TABLE III Rates and Activation Parameters for the Oxidation by $208
k(l mol-lsec -I)
Complex Fe(phen)32+
Fe(bipy)3
2+
Fe(terpy)22+
0s(bipy)3
2+
Ea(keal mol -I)
2-
~S~(cal mol-ldeg -I) Ref
0.108 (25 ° )
13.5
-19
8
0.41 (25 °)
10.3
-28
8
0.85 (35 ° )
10.7
-24
8
49.0 (25 ° )
VO 2+
9
a
ii
+
a.
Reaction only measurable in the presence of Ag .
after which time the reaction started and accelerated with time.
This
behavior suggests that a small concentration of an Os(lll) product is _oi:~=u which is able to catalyze the r~action in the absence of EDTA. Ag + and Cu 2+ catalyzed the reaction strongly.
Even [Cu2+] of i0 -~ M had
While the reaction between Os(phen)32+ and S2082- has
a measurable effect.
not been examined it is expected to be similar to the reaction between
Os(bipy)32+
and $208
2-
~ith a similar rate constant.
Prel~minary reactions of P208
4-
with [Os terpy bipy CI] + suggest that
the rate of hydrolysis of the monodentate CI- ion an~ the rate of oxidation with [P2084-] are similar.
The oxidation reaction is strongly catalyzed by
trace metal ions and the rates are only reproducible in the presence of EDTA. 2A ha!f-life of tl/2 50 hr was observed at pH 7, 60°- Oxidation ~y S208 is extremely fast even at room temperature. Attempts at studying the oxidation of [Ru(phen)2PyCl] + and
6
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. 1
[Ru(phen)2pyC~CN] 2+ by P2084- led only to indications of behavior.
The
reaction of the chloro complex was complicated by hydrolysis of reactant and product.
However, it was established that the complex reacts with
P208 ~- (tl/2 % 1.5 hr at pH 9, 60°) at a faster rate than hydrolysis of #
the con~lex (tl/2 ~ 40 P~ at pH 9, 60o) •
For [Ru(phen)2pyCKjCN] 2+,
hydrolysis is too slo~v to interfere with the rate of oxidation.
However,
the solutions turned out to be photosensitive and half-lives of oxidation varied from
tl/2=
0.5 - 18 hr.
The complications found for Ru(II)oxida-
i]ions are too severe to justify firm mechanistic conclusions. While the reaction of vanadyl ion and peroxodisulfate requires the presence of a silver ion catalyst to proceed at room temperature peroxodiphosphate
(ll) the
reacts at a convenient rate with the vanady! ion.
The
kinetics are flrst-order each in oxidant and reductant with a rate constant of 5 x 10 -2 ! mo! -I sec -I at 25 ° , at pH I.i to 3.5. (12) These results, the different kinetics, the great difference in rate constants, and the very different activation parameters,
clearly demonstrate
that the isoelectronic $2082- and P2084- react with these complexes by quite different mechanisms.
While S208
2-
apparcnt!yprefers
an outer-sphere
mechanism the results from the P2084- reactions (except with the R~(II) complexes) are best explained by an inner-sphere mechanism where" the ratedetermining step is the partial dissociation of a ligand with P2084- or 2PO 4 approaching the vacant position follo~Ted by a fast oxidation reaction. In the Os(phen)32+ and Os(bipy)32+ cases no measurable dissociation takes place and so no oxidation is observed as long as catalysts are absent. Tne oxidation potential for the peroxodisulfate co,~le is -2.01 v. (13) The oxidation potential for the peroxodiphosphate has not been determined but the fact that P2084- oxidizes Ag +, Cu 2+ ~nd V02+ indi.cates that the potential is not ~ich different from that of the $208~C coup] e. It is of interest to esti~ate the minimum activation ener~q[ necessary
PEROXO~PHOSPHATE IoN OXIDATION
VoL 7, No. 1
to explain the marked difference in rates of oxidation of Os(phen)32+ by the two isoelectronic peroxides. difference,
Roughly correcting for the temperature
it is found that $2082- oxidizes the osmium complex at least
108 more rapidly than does P2084-. ii kcal mole -1 in E a.
This is equivalent to a difference of
This seems too larae to attribute to a bond disso-
ciation energy in the peroxide bond. explanation.
There is one attractive possible
If the peroxodiphosphate
is considerably more solvated, thel~
desolvation prior to insertion between the ligand planes (as appears to be the case with peroxodisulfate as oxidant) would be required.
The energy
difference could be rationalized as due to about three hydrogen bonds.
Acknowledg/ments.
We thank the USAF Office of Scientific Research for
continuing support of our research.
i.
Sr. A. A. Green, J. O. Edwards and P. Jones, Inorg. Chem., ~, 1858 (1966).
2.
For the purposes of clarity, peroxodiphosphate will be written as P208
irrespective of protonation state.
This does not alter the basic conclusions.
The state of protonation as a function of pH is reported elsewhere. 3.
M.M.
(3)
Crutchfield and J. O. Edwards, J. Amer. Chem. Soc., 82, 3533
(196o). 4.
F. Basolo, J. C. Hayes and H. M. Neumarm, J. Amer. Chem. Soc., 763
3807 (1954). 5.
j. I~irgess and R. H. Prince, J. Chem. Sot., 6C61 (1965).
6.
R. Kogg and R. O. Wilkin.;, J. Ch~m. Soc., 341 (1962).
7.
R. Farina, R. Hogg and R. G. Wilkins, Inorg. Chem., 7, 170 (1968).
8.
J. Burgess and R. H. Prince,
9-
D. H. Irvine, J. Chem. Soc., 2977 (1959).
iO.
J. Chem. Scc., (A), 1772 (1966).
S. Raman and C. H. Brubaker, Jr., J. Inorg. Nucl. Chem., 31, i091 (1969)
8
PEROXODIPHOSPHATE ION OXIDATION
Vol. 7, No. I
Ii.
D. M. Yost and W. H. Claussen, J..~mer. Chem. Soc., 53, 3349 (1931).
13.
W. H. Latimer, "Oxidation Potentials", Prentice-Hall, Inc., Englewood
Cliffs, New Jersey, 2nd Edition (1952), page 78. 12.
M. Andersen, J. 0. Edwards, Sr. A. A. Green, and Sr. M. D. Wis~ell,
Inorg. Chim. Acta, 3, 655 (1969).