Base Hydrolysis of Chloropentamminecobalt(III) Perchlorat · PDF fileBase Hydrolysis of...
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Base Hydrolysis of Chloropentamminecobalt(III) Perchlorat
and
Visible Spectra of Some Cobalt(III) Comnlexes
by
CHAN Suk-yee
A thesis submitted in partial fulfilment of
the requirements for the degree of
Master of Philosophy in
The Chinese University of Hong Kong
1977
Thesis Committee:
Dr. W.K. Li, Chairman
Dr. K.Y. Hui
Dr. T.C.W. Mak
External Examiner:
Professor Y.T. Lee
(The university of California. Berkeley)
( 陳 淑 宜 )
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2
Acknowledgements
The author wishes to express her sincere thanks to Dr. W.K. Li'
for his guidance and encouragement during the course of her research
and the preparation of this thesis.
She is greatly indebted to Dr. K. Y. Hui and Dr. O. W. Lau for
their discussions on the study of the kinetics.
Thanks are also clue to Mr. L.F. Book for his help at the begin-
ping of this work and to Miss Grace Poon for her skillful tvin.
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Abstract
Part One
Base hydrolysis of chloropentamminecobalt(III) complex was studied
over a wide range of alkali concentrations with both titrimetric-and.
spe trophotometric methods. Excess base was used so that a pseudo-first-
order rate constant was obtained for each run. The rate constants were
first-order in alkali concentration at both 25°C and 0°C. These results
were compared with those obtained by other workers.
Part Two
Visible spectra of some cobalt (III)' complexes were recorded and
interpreted with a crystal field model proposed by Wentworth and Piper.
The overlapping bands of d-d transitions were resolved by the least
squares technique. The crystal field parameters Dq, Dt, Ds,and B were
calculated and compared with those obtained by other workers.
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4
Contents
Part One: Base Hydrolysis of Chloropentamminecobalt(III) Per-
chlorate
I. Introduction1
II. Experimental 10
III. Results and Discussion 13
References22
Part Two: Visible Spectra of Some Cobalt(III) Complexes
I. Introduction24
II. Theory26
ITI. Experimental31
IV. Results and Discussion36
References49
Page
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1
PAR T 0 N E
Base Hydrolysis of Chloropentamminecobalt(III) Perchlorate
I. Introduction
Base hydrolysis, such as represented stoichiochemically by
(1)
has been studied extensively in the past twenty years. However, the
mechanism of these reactions is still unsettled. The rate of base
hydrolysis is about 108 times of that of the acid hydrolysis of the same
complex these reactions obey the rate law
(2)
Such marked enchancement and the clear-cut second-order kinetics led some
workers1 to suggest a simple bimolecular substitution at the central
metal atom.
Study of the effect of added anions to the hydrolysis of
was made by Garrick. It was found that the rate of hydrolysis was
affected only by hydroxide ion. Even the presence of potentially strong
nucleophiles could not affect the rate of hydrolysis. Recognizing that
reactions with similar bimolecular mechanism do not have such an effect,
an alternative mechanism was suggested to explain the unique role played
by the hydroxide ion as the lyate ion (the ion of the solvent).
Garrick' suggested'that the hydroxide ion might act as a base and
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2
served to remove a proton from the amine ligand. Applying this idea
Basolo and Pearson4 developed a conjugate base mechanism which appeared
plausible to account for the extraordinary effect of the hydroxide ion.
The overall process can be represented as:
(3)
(4)
(5)
This mechanism was named the SN1CB mechanism (substitution,
nucleophilic, unimolecular,conjugate base). It involes a fast pre-
equilibrium in which a conjugate base or amido complex is formed so that
hydroxide is the only base to be an effective catalyst. This mechanism
is strongly supported by the fact that complexes lacking an avilable proton
have never been found to be markedly sensitiveto base hydrolysis5. In
these reactions the deprotonated species undergoes a rate-determining
dissociation and eventually forms the hydroxo complex.
If a and b are initial concentrations of complex and-hydroxide,
respectively, and x is the concentration of the conjugate base,
the equilibrium concentrations of complex and hydroxide
are (a-x), (b-x) respectively, and
Since b is much greater than a, the exnressinn may hp gimnlifiPd rn
K
K
fast
K = x/(a-x)(b-x)
K = x/(a-x)b,
X = Kab/(1+Kb).or
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3
The kinetics expression would. then become
Rate= kx
= Kkab/(1+Kb) (6)
When K[OH]« 1, equation (6) is reduced to the observed second-order
rate form, i.e., Rate= Kkab
(7)
where Ka and Kw are the dissociation constant of the conjugate acid and
the ionization product of water respectively.
On the other hand, Chan6 suggested that the reaction was a
bimolecular rearrangement between the coordination shell and the solvation
shell, involing deprotonation of the solvent shell by the hydroxide ion
and Grottus chain transfer to the metal ion. Later, Chan proposed an ion-
pair mechanism7 in which there was a pre-equilibrium formation of an 1:1
ion-pair between the complex cation and hydroxide ion, followed by a rate-
determinating rearrangement within the ion-pair. The rearranged ion-pair
then dissociated rapidly and the hydroxo complex was formed. The overall
process can be represented as:
(8)
(9)
(l0)
The mechanism has been named SN2IP (substitution, nucleophilic,
bimolecular, ion-pair). When the hydroxide ion is in excess, the kinetics
expression is similar to that of the SN1CB mechanism:
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4
(11)
In order to agree with the observed second-order kinetics, the pre-
equilibrium must proceed only to a small extent so that the denominator
in equation (11) approaches to one. That ion-pair association is slight
was satisfactorily explained by the Bjerrum model8 and was supported by
9indirect evidence
Among the various complexes of the form
was known very early and its base hydrolysis was extensively investigated.
The reaction is very fast so that most of the studies were made with low
alkali concentrations. The reaction was found to be second-order overall,
first-order in complex concentration and first-order in alkali concen-
tration.
Adamson et al. 1U studied the kinetics at different temperatures and
interpreted the observed second-order kinetics on the basis of conjugate
base mechanism. Chan 11 studied the base hydrolysis of halogenopentammine-
cobalt(III) complexes at different ionic strengths in aqueous solutions
and concluded that the complexes were expected to follow the pattern of
saturated aliphatic SN2 reactions, viz. FC1(Br or I. Afterwards, Chan12
studied the base hydrolysis of titrimetrically over a
wide range of alkali concentrations at both 25°C and 0°C. The kinetics
was done with excess base at a constant-ionic strength. The pseudo-
first-order rate constants obtained for all runs were shown in Table 1.
The rate constants were found to vary non-linearly with hydroxide con-
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5
centration,as shown in Figure 1. Chan suggested that the departure from
first-order behavior was agreeable with the ion-pair mechanism. Since
the value of the denominator in equation (11) increases with the hydro-
xide-.concentration, a linear relationship should be obtained by plotting
1/kobs against 1/(OH) as shown in Figure 2. This is because equation
(11) can be rearranged into the form:
(12)
Table 1. Pseudo first-order rate constants for the base hydrolysis of
chloropentamininecobalt(III) perchlorate in aqueous solution
12at an ionic strength of 0.1 M.
0.00 0.00 0.00
0.02 0.94 1.25
0.03 1.36 1.79
0.04 1.76 2.32
0.05 2.12 2.80
0.06 2.47 3.25
0.07 2.80 3.68
0.08 3.11 4.08
13
Buckingham et al, repeated the same experiment; however, they
were unable to obtain the same results. The rate of the reaction was
followed spectrophotometrically at 25.4°C. The observed rate constatlts
showed a first-order dependence or. hydroxide concentration (Figure 3).
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Figure 1. Dependence of observed first-order rate constant on [OH] at
25°C and ionic strength 0.1 M for the base hydrolysis of
3
2
1
0
2 4 6 80
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7
Figure 2. Dependence of 1/kobs on 1/[OH] at 25°C and ionic strength 0.1M
for the base hydrolysis of [Co(NH3)5Cl](ClO4)2. 212
80
40
0
0 20 40
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8
Figure 3. Dependence of observed. first-order rate constant on [OH] at
25.4°C and ionic strength O.1M for the base hydrolysis of
[Co(NH3)5C1](C104)2.13
6
5
4
3
2
1
0
0 5 10
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9
Hoping to resolve this difficulty, in this work, the base hydrolysis
of was carried out at O°C and 25°C using both titrimetric
and spectrophotometric techniques.
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10
II. Experimental
1. Preparation
Chloropentamminecobalt(III) chloride was prepared by stand-
ard method.14 The crude product was recrystallized as described
below. Chloropentamminecobalt(III) chloride(66g) was dissolved in
66' ml of concentrated ammonia and 1300 ml of water at 90°C. The
solution-was hot filtered and 132 ml of concentrated hydrochloric
acid was added to the filtrate. The solution was then heated on
steam bath for two hours and cooled. The purified chloride was con-
verted into perchlorate by dissolving in 3M sulphuric acid and adding
70% ice-cold perchloric acid. The crude product was purified by dis-
solving in ice-cold water and adding 70% ice-cold perchloric acid,
and dried at 110°C.
2. Analytical Procedure
An appropriate amount of the complex was treated with excess
of sodium hydroxide. After acidification, the solution was passed
through a column of cation-exchange resin (Amberlite IR-120; H+ form).
The amount of chloride ion in the effluent and washings was determined
by the Volhard method. Found: co-ord.C1, 9.3; Calcd. for [Co(NH3)5Cl]
(C104)2; co-ord.C1, 9.4%
3. Kinetics
Since both reactants are charged the reaction is sensitive
to ionic strength, which was kept constant by the addition of NaC1O4
At O°C the kinetics was studied titrimetrically and spectrophoto-
metrically at ionic strength of O.1M. An appropriate amount of the
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11
cobalt complex was dissolved in NaC104 solution at O°C such that the
complex concentration was less than one tenth of that of hydroxide
ion. Similarly, the sodium hydroxide solution was prepared in water
or in the NaC1O4 solution previously brought to O°C. A big Dewar
flask with crushed ice and water was used as a thermostat. Care was
taken to reduce possible photo-reaction of the complex in solution.
a. Titrimetric Method
Since the reaction is very slow at 00C, ordinary sampling
technique was used. Each time 25 ml of reaction mixture
was withdrawn and the reaction was quenched with ice-cold
perchloric acid. The killed solution was passed through
a column of cation exchange resin (Amberlite IR-120 H+ form).
The column was 15 mm in diameter and 80 mm in length. It
was surrounded with a jacket containing an ice-salt mixture
in order to reduce the possibility of aquation in the
solution when it was passing through the column. The amount
of hydrochloric acid in the effluent and washings was
determined by the Volhard method. A pseudo-first-order rate
constant was obtained by plotting ln(V-Vt) against t where
Vt.and V were the volumes of standard silver nitrate solution
consumed when the reaction was stopped at time t and after
ten half-lives, respectively.
b. Spectrophotometric Method
Without passing through the column, the absorbance at
275 nm of the killed solution was measured immediately
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12
with a Hitachi Model 323 UV-VIS-NIR recording spectro-
photometer. A pseudo-first-order rate constant for each
run was obtained by plotting ln(At- A) against t where
At and A. were the absorbance of the reactant at time t and
after ten half-lives, respectively.
It was unable to study the kinetics at 250C titrimetrically
because fast sampling apparatus was not available. Thus the kinetics
was followed spectrophotometrically only. The reactants were mixed
and placed in the cell which was in the thermostated cell holder.
.The change in absorbance was followed using Hitachi Model 323 UV-VIS-
NIR recording spectrophotometer. The pseudo-first-order rate constants
were obtained as described above.
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13 v-v -«•! f- r« «-» HT34 n m1 r» n i rv n
At 25°C the kinetics was followed spectrophotometrically at ionic
streneth The results are given in Table 2. The
second-order rate constant k is calculated from the expression
The last column of the Table gives the results obtained
o 13at 25.4 C by Buckingham ert al. Under the present conditions the results
12in this work are different from those of Chan but are in good agreemer
13with Buckingham's. The pseudo-first-order rate constants show a first
order dependence on (OH j. Linear plots of the observed rate constants
against are shown in Figures 4 and 5
resDectivelv.
Table 2. Rate Constants for the Base Hydrolysis oi
0.010
0.025
0.050
0.100
0.536
1.32
2.63
5.37
: 0.01
: 0.01
: 0.19
b 0.24
0.54
0.53
0.53
0.54
0.62
0.61
0.55
0.010
0.025
0.050
0.100
0.231
0.57
1.08
2.26
: 0.011
: 0.032
- 0.05
: 0.12
0.24
0.23
0.22
0.23
0.26
0.25
c From Reference 13.
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Figure 4. Rate constants for the base hydrolysis of
measured by spectrophotometric method at 25°C and u = 0.1 M.
6
5
4
3
2
1
00 5
10
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15
Figure 5. Rate constants for the base hydrolysis of
(Co(NH3)5C1)(C104)2 measured by spectro-
photometric method at 25.0°C and µ = 1.0 M.
2
1
0
0 5 10
102 [OH-)(M)
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At 0°C the kinetics was studied spectrophotometrically and titri-
metrically at n = 0.1M;the results are shown in Table 3 and Table 4,
respectively. The observed pseudo-first-order rate constants were plotted
against (oH ] as shown in Figures 6 and 7 respectively. Almost identicalII
resiilts were obtained by both methods, as should be the case. In contrast
to Chan results, the rate constants measured by both methods show no
deviation from first-order dependence on £oH J.
Table 3. Rate Constants for the Base Hydrolysis of
at 0°C Measured by Spectrophotometric Method
0.03
0.04
0.05
0.06
0.08
2.76
3.73
4.81
5.54
7.59
0.02
0.06
0.07
0.08
C. 18
9.20
9.33
9.62
9.23
9.49
Tablp 4. Rate Constants for the Base Hydrolysis of
at 0°C Measured by Titrimetric Methoc
0.03
0.04
0.05
0.06
0.08
2.85
3.71
4.68
5.57
7.72
0.11
0.16
0.15
0.14
0.10
9.50
9.28
9.36
9.28
•9.65
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Figure 6. Rate constants for the base hydrolysis of
measured by spectro¬
photometry method at and
8
6
4
2
0
0 2 4 6 8
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Figure 7. Rate constants for the base hydrolysis of
measured by titrimetric
method at and the present
results from reference 12.
8
6
4
0 2 4 6 8
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Monk studied the base hydrolysis of chloropentamminecobalt(III)
o 36perchlorate at 0 C by following the rate of release of CI at relative-
ly high hydroxide concentrations. The second-order rate constants at
zero ionic strength, k0, were calculated from the second-order rate
constants, k(exptl.)1s, at various alkali concentrations and, therefore,
ionic strengths. The constancy of k0 implies that k at constant ionic
strength does not decrease with increasing (OH J. Thus the pseudo-first-
order rate constants are first-order with respect to OH J. The second-
order rate constant at foH 1 = 0.07960M and u = 0.1079M is
which is in fair agreement with the present result.
12Chan claimed that the substantial deviation from first-order de¬
pendence on OH j of the rate constants was due to the ion-pair formation..
He also studied the base hydrolysis of cis-chloroamminebis(ethylenediamine)-
cobalt(III) and cis-chloroamminetriethylenetetraminecobalt(III) perchlorates
at high alkali concentrations. The plots of the observed rate constants
against alkali concentrations for these cation also showed departures from
the first-order behavior similar to that shown in Figure 1 except that
the curvatures were less marked. It was found that the reaction rates
decreased when ammonia molecules were replaced by multidentate amines. One
of the effects of an increase in chelation is to increase the acidity of
the N-H bonds, the K values. Another effect is to increase the size ofa
the complex so that the ion-pair association constant, K. , is decreased.IP
Consequently a conjugate base mechanism predicts an increase in chelation
would be accompanied by an increase in reaction rate. On the contrary
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the ion-pair mechanism predicts the reverse which was supported by
Chan's results. However, like other workers, we were unable to re-
produce his results.
It is noteworthy that the conjugate base mechanism can also lead
to a departure from second-order kinetics. Provided that the alkali
concentration is sufficently high, a stage will be reached such that
all the original complex is converted into the conjugate base. In
this case further increase in alkali concentration would be no longer
accompanied by an increase in reaction rate. It is also noteworthy
that the present results will be consistent with the ion-pair mechanism
proposed by Chan if the ion-pair association constant is small, e.g.,
A comparison of two mechanisms indicated that the rate expressions
are similar except that the constants k, K and K, have different mean-a ip
ings: k is defined as the rate constant of dissociation of CI from the
conjugate base in the SICB mechanism but as the rate constant of re¬
arrangement within the ion-pair in the S 2IP mechanism; K is the dis-N a
sociation conatant of the conjugate acid-in the former mechanism, while
16K. is the ion-pair association constant in the latter. In 1969 Chan
1P
suggested that the ion-pair and the conjugate base mechanisms are
similar and differ only in the extent to which the proton is transfer
from the complex to the base. If the acidity of the proton is high, it
will stay close to hydroxide ion and this situation approximates to a%
conjugate base On the other hand, if the acidity of the proton is low,
it will stay close to the complex and this situation then approximates to
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an ion-pair. Unfortunately nothing is known about the acidity of
chloropentamminecobalt(III) complex except that it is extremely low,
nor the ion-pair association constant, K . Chloropentamminecobalt(IIIip
complex appears to be too weak an acid for direct measurement of
it ionization constants. Direct measurement of K is also extremely
difficult because the association is slight and because the cation is
too reactive towards OH .
• Although most of the experimental results can be equally well ex¬
plained by both mechanisms in the past, recently some workers'
claimed that the conjugate base mechanism was closer to the truth.
Based on the constancy of the enthalpy change from transition state to
products of the series of halide complexes it was sug¬
gested that the reaction mechanism was dissociative and the S2IP
mechanism was excluded. However, an SlIP mechanism can occur in
these reactions. In addition, spectrophotometric evidence for the
formation of the conjugate base rather than of the ion-
pair in strongly alkali solutions of
18
provides further support for the SICB mechanism . However,
is a weaker acid than and the formation of
rather than is suspect in the author's opinion.
Based on the results of this work, a preference of one mechanism
over the other cannot be made. We can only re-emphasize that the base
hydrolysis of obeys a second-order kinetics perfectly at
high alkali concentrations.
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References of Part One
1. (a) C.K. Ingold, R.S. Nyholm, and M.L. Tobe, J. Chem. Soc. , 1691(1956).
(b) C.K. Ingold, R.S. Nyholm, and M. L. Tobe, Nature (London), 194,
344(1962).
(c) R.S. Nyholm,and M.L Tobe, J. Chem.Soc. , 1707(1956).
(d) D.D. Brown, C.K. Ingold, and R.S. Nyholm, ibid., 2674(1953).
(e) D.D. Brown and C.K. Ingold, •ibid.. 2680(1953).
2. (a) F.J. Garrick, Trans. Faradav Soc., 33, 486(1937).
(b) F.J. Garrick, ibid., 34, 1088(1938).
3. F.J. Garrick, Nature. 139. 507(1937).
4. F. Basolo and R.G. Pearson, Mechanisms of Inorganic Reactions,
2nd ed., Wiley, New York, N.Y., 1967, pp. 173-193 and 261-265.
5. M.L. Tobe, Acc. of chem. Res., 3, 377(1970).
6 S.C. Chan and M L. Tobe, J. Chem. Soc., 4531(1962)
7. S.C. Chan and F. Leh, ibid. (A), 126(1966).
8. N. Bjerrum, Kgl. Danske Videnskab Selskab Mat.-fys. Medd., 9, 7(1926).
9. A.W.'Adamson and R.G. Wilkins, J, Amer. Chem. Soc., 76, 3379(1954).
10. A.W. Adamson and F. Basolo, Acta Chem. Scand., 9, 1261(1955).
11. S.C. Chan, K.Y. Hui, J. Miller, and W.S. Tsang, J. Chem. Soc., 3207(1965).
12. S.C. Chan, ibid.(A), 1124(1966).
13. D.A. Buckingham, I.I. Olsen, and A.M. Sargeson. Inorg. Chem., 7,
174(1968).
14, H.H. Willard and D. Hall, J. Amer. Chem. Soc., 44, 2220(1922).
15. M.R. Wendt and C.B. Monk, J. Chem. Soc. (A). 1624(1969).
16. S.C. Chan and O.W. Lau, Aust. J. Chem., 22, 1851(1969).
17. D.A_ House and H.K.J. Powell, Inorg. Chem., 10, 469(1971).
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18. L. Heck. Inorc. Nucl. Chem. Lett., 6, 657(1970).
19. J. Burges(Senior Reporters), Inorganic Reaction Mechanisms, Vol. 1,.
'A Review of the Literature published between January 1969 and August
1970', A Specialist Periodical Report, The Chemical Society, Burlington
House, London, 1971, pp. 177-182.
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PART TWO
Visible Spectra of Some Cobalt(III) Complexes
I. Introduction
The visible spectra of monoacidopentaammine and trans-diacido-
tetraammine trivalent cobalt complexes have been studied extensively.
The splitting of the first two spin allowed bands that occurs upon the
lowering of the symmetry from 0 to and has been the subject
of numerous investigations, both theoretical and experimental. Inter¬
pretations and predictions in terms of a crystal field model were given
2 3.4by Moffitt and Ballhausen and Yamatera , while McClure and also Yama-
3tera used approaches based on a molecular orbital model.
Regardless of the model used, the splitting of both bands must be
known to evaluate the crystal field parameters. In general, as pointed
out by Linhard and Weigel in 1951, the splitting of the first band;
which is lower in energy than the second, will be readily observed if
the axial ligand is well separated from ammonia in the spectrocnemical
series. On the other hand only a broadening or band shift will be ob¬
served if the separation is not large. In addition, the splitting of
the high energy band, the second band, has never been observed. Thus
evaluation of all the parameters which arise from the theories is ham¬
pered by the fact that the splitting of both bands is required.
In 1965 Wentworth and Piper' recorded the visible spectra of some
monoacidopentaammine, trans-diacidotetraammine, and trans-diacidobis-
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(ethyleriediamine) cobalt (III) complexes and successfully correlated
band splitting with various properties of the ligands. They recog¬
nized that the energies of the first bands in and complexes
are almost exactly that of the transition in the parent
octahedral complex. or the pseudo-octahedral complex,
Thus the splitting of the lowest lying excited state
was resolved in the electronic spectrum. Later other workers'
also applied the Wentworth-Piper model to the interpretation of
and complexes successfully. However the second band was still
unresolved.
gThe most recent and successful effort was that of Book et al.
The visible snert~re nf and with
Et, n-Pr and n-Bu were studied. Since the parent complexes
have not been prepared, the resolution of the bands was
achieved by least squares assuming that the bands were Gaussians.
In this work, the Wentworth-Piper model was used to interpret the
visible spectra of some cobalt(III) complexes but Book's method was
used to resolve the overlapping peaks of d-d transitions.
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II. Theory
The electronic spectrum of an octahedral metal complex with low-
spin configuration is characterized by two bands in the visible
and ultraviolet which may be assigned, in ascending energy order, to
the and transitions respectively. Within
the formalism of crystal field theory, the energies of the lowest-
lying singlet and triplet states above the ground state are
(13)
(14)
(15)
(16)
where B and C are Racah interelectronic repulsion parameters. The
wavenumbers of.the spin-allowed singlet bands are generally used to
solve for Dq with the assumption that the ratio BC has the free-ion
value. Another method is to use the wavenumbers of thesand
bands instead of that of bands and bands.
On descent in symmetry to or corresponding to a tetra-
gonal distortion, the degeneracy of the excited states is partially
removed as shown in Figure 8. The state of 0 splits into
and of and becomes and Follow-
1 laing Wentworth and Piper, the lower E state is designated as E and the
4
the higher as E. To a first-order approximation the energies of the
singlets above the ground state are
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(17)
(18)
(19)
(20)
Figure 8. Tetragonal splitting of the excited states of cobalt(III)
The parameters Dt and Ds were defined in D71 field as4h
and in the C, field as4v
(21)
(22)
(23)
(24)
9The radial crystal field parameters were defined as
(25)
where e and r are the'electronic charge and nucleus-electron distance
respectively, while q and R are the effective ligand charge and distance
from the metal, respectively. The subscripts refer to the coordinate
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axes on which the ligancls arc placed. As an approximation Wentworth
and Piper defined Dq so that it depends only on the in-plane
field strength. In other words, when comparing
the fields of and or trans In
jaddition, for complexes the ligands on xy plane and z- axis areA
amines so that%
and therefore
(26)
(27)
(28)
Comparing with those of D, complexes we now have the relationships
(29)
(30)
Furthermore, Wentworth and Piper assumed that the radial crystal field
parameters are characteristic of the ligand itself irrespective of the
particular substituted complex ion in which it is found. Then the para-4
meter Dtfor C and complexes may be expressed as
(31)
(32)
The splitting of the first band is (354)Dt, while that of the
second is 6Ds - (54)Dt. In the case that both.components of each state
cannot be observed, assumptions are made to evaluate Ds and Dt. Went-
worth and Piper defined a useful empirical parameter Dt' by the equation
(33)
which is an approximation to Dt, neglecting off-diagonal elements of
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the perturbation matrix. Quantity W is the energy of IA band and quan¬
tity was taken from the parent octahedral complex. In
order to extend the calculation to those cases in which the IA and lb
bands were not resolved, the observed band maximun was taken to be the
average of the energy of the transition in the parent
complex and that of the IA band. However the parameters B and Ds could
not be evaluated since they did not resolve the second band.
Later Ban and Csaszar overcame this difficuly by assuming that the
second band has never been split so that
(34)
where E(II) was the energy of the unsolved band II. Therefore Ds and
B could be calculated from the expressions
(35)
and
In the present work the bands are resolved into Gaussian curves by
an iterative least-squares procedure assuming a sloping baseline. The
computer programme was originally written by Frasen and Suzuki. The
xyparameters Dq , Dt, i, Da are calculated from the relations
(37)
(38)
(39)
(40)
z+Dq of a complex is evaluated from the expression
(41)
(42)and
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for a complex. It is assumed that C is a constant irrespective of
substitution of the parent octahedral complex. This assumption is
justified by the fact that C is remarkably constant over a wide ranee
of field strength. Foi thew ~
values of C are found to be 3825, 3835 and 3650 cm respectively and
hence C is taken to be 3.8 kK in this work.
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III. Experimental
1. Preparations
a. Carbonatopentamminecobalt(III) Nitrate
A solution of 30 g of cobalt(II) nitrate hexahydrate
in 15 ml of water was thoroughly mixed with a solution of
45 g of ammonium carbonate in 45 ml of water and 75 ml of
concentrated aqueous ammonia (sp. 0.885, 33% NH ). A
stream of air was bubbled slowly through the mixture for
24 hours. After the mixture had been cooled in an ice-
salt bath overnight, the product was collected on a filter,
washed with not more than 5 ml of ice-cold water, followed
by alcohol and ether, and dried at 50°C. Anal. Calcd.
for Co(NlI3)5C03jN03.%H20: C, A.37; H, 5.86; N, 30.55;
Found: C, 4.28; H, 5.96; N, 30.84.
12b. Aquopentamminecobalt(III) Oxalate
Chloropentamminecobalt(III) chloride was synthesized
and recrystallized as described in Part One. A mixture
of 10 g of finely powdered Co (NH) Cljc, 75 ml of water,
and 50 ml of 10% ammonia was heated on a steam bath in an
Erlenmeyer flask covered with a watch glass with continuous
agitation until all of the basic aquopentammine chloride
dissolved and a deep-red solution formed. The solution was
filtered, the filtrate was made very weakly acidic withi
oxalic acid, and some additional ammonium oxalate was added
to complete the precipitation. The slurry was allowed to
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stand; the precipitatewas then filtered off and washed
with cold water. Anal. Calcd. for
(c2°4)3AH2°: c 10.91; H, 6.41; N, 21.21; Found: C,
10.78; II, 6.15; N, 20.96.
13c. Nitropentamminecobalt(III) Chloride
A mixture of 20 g of 200 ml of
water, and 50 ml of 10% ammonia was heated with successive
shaking until the salt dissolved. The solution was filt¬
ered. The filtrate was cooled and made weakly acidic with
dilute hydrochloric acid. About 25 g of crystalline
sodium nitrite was added and heating on the steam bath was
continued until the initial red precipitate dissolved
completely. The cold, brownish-yellow solution contained
a copious deposit of crystals. Then 250 ml of concen¬
trated hydrochloric acid was added. After chilling, the
product was filtered, washed with 1:1 hydrochloric acid,
and then with alcohol until free of acid, and dried in
air. Anal. Calcd. for H, 5.79; N,,32.20;
Found: H, 5.87; N, 32.60.
14d. Propionatopentamminecobalt(III) Nitrate
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Carbonatopentamminecobalt(III) nitrate (5 g) wast
suspended in 15 ml of water, and 15 ml of n-propionic
acid was added. The reaction mixture was concentrated
on a steam bath for 1.5 hours, during which time a red
crystalline was separated. After cooling to room tem¬
perature, 50 ml of water was added. The product was
filtered and washed wi.th 50 ml of cold water, followed
by alcohol and ether, and dried at 50°C. Anal. Calcd.
for C, 7.34; H, 5.55; N, 29.97;
Found: C, 7.27; H, 5.37; N, 30.10.
e. Fluoropentamminecobalt(III) Nitrate
This compound was kindly supplied by Dr. K.Y. Hui.
It was recrystallized by dissolving it in minimun amount
of water at 45°C and adding ammonium nitrate to the
solution after cooling to 0°C. The purified product was
collected and dried at 90°C.
f. Nitratopentamminecobalt(III) Nitrate
Carbonatopentamminecobalt(III) nitrate (10 g) was
suspended in 25 ml of water, and 20 ml of nitric acid
(1:1 concentrated acid and water) was added with stirring.
When the evolution of carbon dioxide had ceased, 100 ml
of methanol was added, the aquopentamminecobalt(III)
nitrate was filtered, and then washed with alcohol
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and ether. This salt was heated at 100 C until 1 mole
of water was lost (about 18 hours). Anal. Calcd. for
H, A.58; N, 33.95; Found: H, A.75
N, 33.83
g. trans-Dichlorobis(ethylenediamine)cobalt(III)
Chloride
Sixty grams of a 10% solution of ethylenediamine was
added, with stirring, to a solution of 16 g of cobalt(II)
chloride hexahydrate in 50 ml of water. A vigorous
stream of air was drawn through the solution for 8 hours
Then 35 ml of concentrated hydrochloric acid was added ai
the solution was evaporated on the steam bath until a
crust formed over the surface. The solution was allowed
to cool and stand overnight. The bright green square
plates of the hydrochloride of the trans form were filt¬
ered, washed with alcohol and ether and then dried at
110°C. At this temperature, the hydrochloride was lost
and the crystals turned into dull-green powder. Anal.
Calcd. for trans
Found
18h. trans-Dibromobis(ethylenediamine)cobalt(III) Bromide
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Ten grams of trans-dichlorobis(ethylenediamine)
ccbalt(III) chloride was dissolved in 50 ml of concen¬
trated hydrobromic acid. The reaction mixture was
evaporated to dryness on a water bath. Then 50 ml of
Concentrated hydrobromic acid was added and the eva¬
poration was repeated. The residue was washed with
.
cold water and then with alcohol. Anal. Calcd. for
trans C, 11.47; H, 3.85; N, 13.43;
Found: C, 11.38; H, 4.18; N, 13.40.
2. Spectral Measurements
The ultraviolet and visible spectra of the complexes were
recorded with a Hitachi Model 323 UV-VIS-NIR recording spectro¬
photometer. All measurements were made at room temperature.
The ultraviolet spectrum was obtained from 210 nm to 340 nm
while the visible spectrum was obtained from 340 nm to 700 nm.
Reagent grade methanol was used as solvent for trans
CI and trans Br because they are known to be un¬
stable towards hydrolysis and isomerization in water. For other
complexes water was used since their aquation rates are extremely
19slow. All the spectra were recorded as soon as the complexes
had been dissolved. The scanning time was about seven minutes
to sweep the whole spectrum. Repeated recordings showed no change
of the spectra.
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IV. Results and Discussion
The electronic spectra and the resolution of the overlapping
peaks of all the complexes are shown in Figures 9-16. The wave¬
lengths intensities and halfwidths of the resolved bands are given
in Table 5 together with those reported by other workers for com¬
parison. The crystal field parameters calculated from the experi¬
mental results given in Table 5 along with those obtained by other
workers are summarized in Table 6 including those of
andg
obtained bv Book et al. The second bandsof
trans and trans-
were not resolved due to the onset of the charge
transfer bands and therefore the values of Ds and B are missing.
J
Upon examining the results, the following remarks can be made:
(1) The order of the values of Dq obtained from crystal
field treatment is NC
Br , an arrangement in good agreement with that
n • j 20normally listed.
(2) Comparing the values of Dq(Cl), Dq(Br) of the trans-
disubstituted complexes with those of the monosubstituted
8complexes reported by Book et al., it is found that these
values remain fairly constants in C. and Dt_ complexes.J 4v 4h
Also the values of Dq(NH), Dq(en) are almost unchanged
irrespective of the particular substituted complex ion in
which it is found. Thus the assumption that the crystal
field strength is characteristic of the ligand itself is
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justified.
(3) The comparison of the parameters Dt and Ds of the and
D., complexes indicates that the values of the latter are about4h
twice as large as those of the former as predicted by the
electrostatic model of the crystal field theory.
(4) In the case where the IA and IB bands were not resolved,
Wentworth and Piper took for granted that the observed band
maximun was the average of the energy of the
j £
transition of the parent compound and that of the E state
above the ground state. This approximation in effect assumed
that the intensities of IA band and IB band were equal. How¬
ever, this assumption is unreasonable since the
transition which corresponds to IB band is symmetry forbidden
in both and symmetries. As seen from the results
given in Table 5 the intensities of IA and IB bands are somewhat0
different.
(5) In the determination of the Dt value, Wentworth and Piper
assumed that the energy of the IB band was approximately equal
to that of the transition in the parent complex,
which was equal to 21.05 kK. As seen from the
results given in Table 5, the energies of the resolved IB band
reported by them and those obtained in this work show no sig-
nificant deviation from this value. However our results of Dt
do not agree with those reported by them. Obviously their
assumption is not always justified as seeii from the data of
reported by them. If the energy of their IB
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band is used in the calculation instead of the first band of
the parent complex, a value close to ours will be obtained,
which is about 26% greater than theirs.
(6) The values of Ds obtained by Ban are very much smaller
than ours as a result of their neglecting the splitting of
the second band. Unfortunately, the crystal field model does
not make a clear-cut prediction of the splitting of the T
3level. Based on a molecular orbital model, Yamatera sug¬
gested that the splitting of band II should be less than that of band
I (2000 cm ) . Since our results do not contradict this pre¬
diction our values of Ds calculated from the splitting of band
II do not seem to be unreasonable. Furthermore the valuer of
Ds of and8
obtained by Book e_t al.
and that of trans-J CI calculated from the polarized
crystal spectrum by Dingle are also substantially greater than
tho.se obtained by Ban.
(7) Unlike other complexes the energy of the A0 state above the
ground state off 1 SL
is less than that of the E state
since the field strength of N0 is greater than that of NH.
In conclusion, it is believed that a rather simple crystal field
model is successful in the interpretation of the electronic spectra.
In addition, the resolution of the bands can lead to consistent, and
therefore helpfully meaningful, values for the crystal field parameters.
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QuCo
_Qv_O00_£
0. 6
0.4
0.2
0.0
Fir.ure 9. Resolution of the shouldered band into IA and IB
transitions foi
solvent:
concentration:
IB
IA
haspli np
350 400 450 500 550 5 00 650
Wa velenq th, n m
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our~
O_o
k_oCO
-Q
o.s
0.6
0.4
0.2
O.C
220 • 360 400 440 400 520 560 6
Wavelength,nm
Fieure 10. Resolution of the shouldered bands into IA, IB, IIA and
IIB transitions for
solvent:
concentration:
IB
IA
n b
n.A
base!rne
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1VJC
j:k.cw,.£
2.4
'1.8
1.2
0.6
0.0?QQ 340 390 440 4 9 O 540
Wavelength, nm
F1pure 1 1 . Resolution of the shouldered bands
into IA. IB. IIA and IIB transitions
for
solvent:
concentration:
IIB
IIA
baselineIR
IA
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cOr
c_c
t-cV
JZ
o.e
0.4
0.2
0.0
F1 Piirp 17 . Resolution of the shoul dered hpnds into TA. TR . TTA end TTR rrenci'rinnc
f OI
solvent:
concentration:
Tt F
2 AIB
T A
hornl !nr
- ' ' 1 I I S J I t L t I LI : I » 1 t- t r J 1 A 1- 1 J I 1 1 I .
320 360 400 44 0 4 8 0 520 560 60 0 640
WavelenQ th. nm
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0uCo
_0k_Otn
_o
o.:
0.2
0.1
3 1«— iii J; i i i i ——J-' L i—--—.320 360 400 4 4 0 480 520 560 600 640
Wavelength ,nm
i
VM HResolution of the shouldered bands into IA_ TR TTA ttr
transitions for
solven t:
concentration:
IA
TT F
n5'ib
base I ine
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G)uCo
JDJ—oI)
_Q
0.3
0. 2
0.1
340 390 440 490 540 590 640
Wavelength, nm
Fi on tp 1 L Resolution of the shouldered band into
IA and IB transitions for
solvent:
concentration:
E
ba se! i npIB
TA
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D'wCo
_Qk_OCO
_a
0.3
0.2
0.1
n.o
Figure 15 . .Resolution of the overlapping peaks into
1A and IB transitions for t
solvent: MeOH
concentration:
II
I B
IA
basel ine
34 n 390 440 490 540 5 90 640 6 90
Wavelenqfh.nm
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o
cC
_Q
O00r
0.6
0.4
0.2
o.o440 490 540 590 640 690
Wavelengnh ,nrn
Figure 16. Resolution of the overlapping peaks into IA
and IB transitions for
solvent: MeOH
concentration:
baseline
IB
IA
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Ihble 5 . Hie wavelengths in kK), intensities (log £.being molar extination coefficient), and half width
of the resolved bands.
Complex Band IA Band IB Band IIA Band 113
13.00 1.7'- 2.01 21.10 1.68 4.13 26.38 (shoulder; net resolved)
l8.90a 20.503 27.35° (not resolved)
19.34 I.63 3.39 21.05 1.83 4.16 27.86 1.09 3.7'+ 29.50 1.79 3.93
20.50° (not resolved)
?0.30b (not resolved)
30.40 (not resolved)
30.30 (not resolved)
23.87 1.51 5.13 21.51 1.29 4.25 31.65 2.55 4.94 28.57 2.66 3.52
21.75° (not resolved)
21.84b (not resolved)
30.80a (not resolved)
30.80b (not resolved)
19.38 1.67 2.87 20.70 1.58 3.5 28.09 1.73 3.1 29.94 1.47 0.90
19.37°
10.88b
21.35°
21.00b
23.41° (not resolved)
28.39 (not resolved)
19.05- 1.97 2.9c
19.45®
19. 45b
21.37 1.73 3.'+5 27.70 1.70 3.62 . 29.07 1.70 3.76
21.80
21.-'»7b
28.27 (not resolved)
28.27k (not resolved)
19.54 1.61 3.16
19.203
21.37 1.45 'f .15n
PI P0
28.90 . (shoulder; not resolved)
28.0° (not resolved)
16.34 1.39 2.53
16.20°
16.12b
16.48°
21.84 1.40 4.43 25.01 (not resolved)
25.85° (not resolved)
25.92b (not resolved)
27.10° (not resolved)22.50° 24.25°
15.24 1.72 2.20
15.25°
15.21b
21.69 1.39 2.93
21.72°
21.68b
Masked
2612°
Maskedb
a. from reference 6 b. from reference 1 c. From reference 21
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Crystal field parameters (in kK) calculated from the spectral data
-l Dt Ds 3
2.490 1.612
1.633'
0.251
o..46a 0-03i3a
n laq n.k
: nc;Aa -n i o£a n no£a n
7.4QOb i.88Sb n.i7Pb
7.331 3.476 _n.P70 -n.mo 0.477
3.038a -o.i 6na -o.0333a o.33ia
a _iiqnb i -ob _n i8ob
2.430 1.922 0.131 0.340 0.394
p n?na n iana n °o3a n c;aoa
o UQP,b a n iUb
2.517
a han
1.590
l.S58a
l_34qb
0.265
0.l83;
0.1351
0.284
0.03Sla
0.59
0.5l8a
3.17 1.705
1 - 738
0.232
1.plf 0.0443 0.536a
o i !i7 n C,~Ci
2.530b
0 £aoc
1.483s
1.459'
1.426'
0.600'
0.612'
9 687 n 6vlc
2.59
2.550'
1.259
1.293aV.
1.777
0.737
0.709'
0.716'
0.l4?a 0.550a
2.51 1.46 0.299 0.227 0.301
2.48 1.31 0.33
xy , 7. •+Dq' denotes the Dq value of NH or on; Dq denotes the Dq value of the
tprni n7 nrr 1 7 :rnnHf Q)
a. From reference 6.
b. From reference 1.
c. From reference 21
d. From reference 8b
e. From reference 8a
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References of Part Two
1. For a very complete set of references, see R.A.D. Uentworth and
T.S. Piper, Inorg. Chem., 4, 709(1965).
2. W.Moffitt and C.J. Ballhausen, J. Inorg. Nucl. Chem., 3, 178(1956).
3. H. Yamatera, Bull. Chem. Soc. Janan. 31. 95(1958).
4. D.S. McClure, Advances in the Chemistry of Coordination Compounds,
S. Kirschner, Ed., The Macmillan Co., New York, N.Y., 1961, p.498.
5. M. Linhard and K. Weigel, Z. anorg. allgem. Chem., 264, 321 (1951) ; ibid. ,
266, 49(1951): ibid., 267, 113(1951); ibid., 267, 121(1951); ibid.,
271, 101(1952).
(a) M. Ban and J. Csaszar. Magv. Kern Foly. 73 (11), 509(1967); ibid.,
73 (11), 512(1967); ibid., 74. (8), 333 (1968) ; ibid., 74_ (12), 587
(1968) .
(b) M. Ban and J. Csaszar. Acta Chimica Academiae Scientiarum
Hunearicae Tomus 57 (2). 153(1968).
7. G.R. Brubaker and J.J. Fitzgerald, J. Coord. Chem., 4, 93(1974).
8. (a) L.F. Book, K.Y. Hui, O.W. Lau, and W-K Li, Z. anorg. allgem. Chem.
426, 215(1976).
(b) L.F. Book, K.Y. Hui, O.W. Lau, and W-K Li, ibid., 426, 227(1976)
9. T.S. Piper and R.L. Carlin, J. Chem. Phys., 33, 1208(1960).
10. R.D.B. Fsaser and E. Suzuki in Spectral Analysis; Methods and Tech¬
niques, J.A. Blackburn (ed.), Maecel Dekker Inc., New York, 1970,
p.171.
11. F. Basolo and R.K. Murmann, Inorganic Synthesis, Vol. 4, H.S. Booth
i
(ed.). McGraw-Hill, 1953, p.171.
12. S,M. Jorgensen, Z. anorg allgem Chem., 19, 78(1899).
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13. G. Brauer, Handbook of Preparative Inorganic Chemistry, 2nd ed.,
Vol. 2, Academic Press Inc., New York, N.Y. 1003, p. 1534.
14. F. Basolo and R.K. Murmann, Inorganic Synthesis, Vol. 4, H.S.
Booth (ed.), McGraw-Hill, 1953, p.175.
15. F. Basolo and R.K. Murmann. ibid.. Vol. 4. d.172.
16. F. Basolo and R.K. Murmann, ibid.. Vol. 4, d.174.
17. John C. Bailar. Jr.. ibid.. Vol. 2. d.222.
18. S.M. Jorgensen. J. Prakt. Chem.. 41 (2). 440(1890).
19. F. Basolo and R.G. Pearson, Mechanisms of Inorganic Reactions,
2nd ed., Wiley, New York, N.Y., 1967, p.164.
20. C.K. Jorgensen, Absorption Spectra and Chemical Bonding in Com¬
plexes, Pergamon Press Ltd., London, 1962, p. 109.
21. -R. Dingle, J. of Chem. Phvs. . 46, 1(1967).
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