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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 )

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.

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.

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

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

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

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:

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-

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).

6

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

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

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

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.

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

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

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.

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.

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

15


(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)

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


measured by spectro¬

photometry method at and

8

6

4

2

0

0 2 4 6 8


measured by titrimetric

method at and the present

results from reference 12.

8

6

4

0 2 4 6 8

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

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

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.

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).

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.

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-

(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.

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

(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

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

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

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.

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

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

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

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

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.

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

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

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.

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

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

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

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

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

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

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

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

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)


23.87 1.51 5.13 21.51 1.29 4.25 31.65 2.55 4.94 28.57 2.66 3.52


21.84b (not resolved)

30.80a (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



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


28.27k (not resolved)

19.54 1.61 3.16

19.203

21.37 1.45 'f .15n

PI P0

28.90 . (shoulder; not resolved)


16.34 1.39 2.53

16.20°

16.12b

16.48°

21.84 1.40 4.43 25.01 (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

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

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).

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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