Thermodynamic properties of actinide complexes—II

J. inor~ nail. Chem.. 1975, Vol. ~7, pp. 2t77 2179. Pergamon Press. Printed in Great Britain

THERMODYNAMIC PROPERTIES OF ACTINIDE COMPLEXES--II THORIUM(IV)-ACETATE

SYSTEM

R. PORTANOVA, P. Di BERNARDO, O. TRAVERSO,* G. A. MAZZOCCHIN and L. MAGON* lsliluto di Chimica General< Universita di Padova and Laboratorio di Chimica e Technologia dei Radioelemenli del C.N.R. Padova, ltal} {Received 5 August 1974) Abstract--The changes in free energy, enthalpy and entropy for the formation of thorium(IV)-acetate complexes, have been determined at 25°C and in an aqueous perchlorate-medium 1.00 M. All five complexes formed, are found to be stabilized by a large gain of entropy. The marked increase in the values of AH, and iS,, relative to the third step of complexation, has been explained in terms of a coordination change. INTRODUCTION WE PREVIOUSLY reported the changes in the thermodynamic functions for the formation, in aqueous solution, of uranyl(VI) complexes with some monocarboxylate ligands[1]. All the complexes were found to be stabilized by a large gain of entropy, while the enthalpy term opposes the complex formation. It is of interest to study how these thermodynamic functions vary when complexes are formed between different actinide ions and the same series of monocarboxylate ligands. We now report a study of the thorium(IV)-acetate system. The stability constants of three successive complexes of thorium(IV) with acetate have previously been determined at 20°C and an ionic strength of 1.00 M made up with NaC104 [2], but no enthalpy or entropy data for this system have been reported. EXPERIMENTAL

The notation and the general calculation procedure were as reported previously[l]. Concentrations are expressed as molarities. As in [1], a temperature of 25°C and an ionic strength of 1.00 M, made up with sodium perchlorate, were used. Chemicals A stock solution of thorium(IV) perchlorate was prepared and standardized as before [2]. Buffer acetate solutions, with a C.dC~,~ (L = acetate) ratio of 4:1 (4 M HL and 1 M NaL) were prepared by parual neutralization of a sample of acid (analytical grade) with standard NaOH solution. Determination of stability constants Corresponding values of [L-] and ti were obtained according to the following equations: [L ] ~ (C~.,_ + C. - [H+])/K"[H ~] and ~i = (CN.L + [H +] - Ca - [L-])/CM. A value of K " = 4.06 x 10" was used [1]. By applying Fronaeus' extrapolation method[3] to the corresponding values of [L-] and h, the overall formation constants,/3j, were calculated for j = 1,2, 3, 4 and 5. The values of/3~ so obtained were then refined by the least squares program "Gauss Z", which minimizes the sum of the squares of the residual (,'L~,- rLb3[4]. *Istituto Chimico di Ferrara, Italy.

Determination of enthalpy changes The AH~ values were first obtained graphically[I]; these were then refined by the least squares program "Letagrop-Kalle" [5.6], which minimizes the error squares sum: ,U : E,W,(Q ............. - Q~.~,,rrf. The weight of each value of Q~..... was equal to unity. From the enthalpy changes and the changes in free energy computed from the corresponding stability constants, the entropy changes were calculated from: AG : AM - TA&. The calorie was taken to be equal to 4.184 absolute joules. Potentiometric measurements The procedure described previously was used[l]. Known volumes of a solution S, (composition: C°,L.,C~D were added to a known volume of a solution S_, (composition C,~ in Th(CIO4)4; C:~ in HC104; (1--4C~-C~) in NaCIO,) in the titration half-cell. Calorimetric measurements An LKB 8721-.2 Precision Calorimeter was used. The heat equivalent, 6., of the calorimeter system was determined by several series of electric calibrations; the standard deviation of the mean value of ~, was 0.05%. The titration procedure was as before[l].

RESULTS

Metal complex [ormation The potentiometric data are reported as a plot of a vs - l o g [L ] in Fig. 1. They show that ~ varies neither with C~ nor with C~; therefore, no polynuclear or acid complexes are formed in the ligand-metal concentration range investigated. Hydrolytic reactions of the thorium(IV) ion are also negligible under the conditions investigated [7]. A maximum ri value of 4.3 was reached, so we were able to determine the stability constants for five successive mononuclear complexes (Table 1). Our previous investigations[2], referring to the same ionic strength and 20°C, led to the determination of only three successive mononuclear complexes; however a less concentrated buffer solution was used. The values of the stability constants previously determined are in good agreement with those reported in Table 1. The calorimetric data for each titration are reported as values of total molar enthalpy changes, Ah~, against

2177

2178

R. PORTANOVAet

al.

Table 1. Stability constants and the values of free energy, entbalpy and entropy for the stepwise reactions of the thorium(IV)-acetatesystem in 1.0 M NaC10, at 25°C. ~j

A }lj

/~ Gj

( M -1)

(Kcal mole -1)

A Sj

(Kcal mole -l)

(cal mole-ldegree -I)

1

(7.24 + 0.30).IO 3

5.26 + 0.03

2.70 + 0.04

26.7 + 0.2

2

(9.39 + 0.45).106

4.u95 + 0.03

1.O7 + 0.03

17.8 + 0.3

3

(8.73 + 0.64).108

2.69 + 0.03

3.27 + 0.09

20.0 + 0.3

4

(1.94 _+ 0.26).iO IO

1.86 + 0.03

1.24 + O.21

10.4 + 0.8

5

(9.90 + 1.50).IO IO

0.97 + 0.03

0.89 + 0.54

6.2 + 1.8

1o

,

M

~o C~i

Thoriu m or)- aeetat •

~'. o zx o

m.

1000

MLz

c'. ~M 2891

~o.eo

28•74

40~

2@24

o(

50

40

30-to0[C] 20

~o

Fig. 3. The distribution of thorium(IV) among the different complexes, MLj, as a function of free ligand concentration. g

s~o

,.o

2o -,0, It]

3.o

,o

Fig. 1. The complex formation curve. The curve was calculated from refined.complexityconstants. values in Fig. 2. The good overlap of the data indicates that the function is not dependent on C~I and C] and this further confirms that polynuclear complexes and hydrolytic reactions are not important. The highest Aho values reported in Fig. 2 correspond to values of about 3.8; here about 15 per cent of the fifth complex is formed (Fig. 3). This allows us to calculate an Thorium(W)

8000

Ahvq 0

acetate

0~ mM

C~ m M

10.00

28.91

$ 10,00

6.96

z~ 20,00

28,74

"o

Z~oo

~o

J

o

2!o

3!o

Fig. 2. The total molar enthalpy change, Aho,as a function of ri. The curve was calculatedfrom the values of/3j and AM.

approximate value of the enthalpy change relative to the fifth step, AHs. The values of free energy, enthalpy and entropy changes for the thorium(IV)-acetate system are reported in Table I. The limits of error refer to standard deviations obtained from the computer programmes. The values of the heats of protonation and dilution of the acetate ligand agree with those determined before[I]. DISCUSSION The curve in Fig. 1 shows an inflection around r~ = 2, due to a fall of the stability of the third complex compared to the second one. This is confirmed by the higher ratio value of the step stability constants K2/K3 = 13.9 compared with Kt/K2 = 5.6 and is clearly shown in Fig. 3, in which the maximum in the ML2 curve is higher than that for ML3. The values of the other ratios are K3/K, = 4.2 and K4/K5= 4.3. The A h : c u r v e (Fig. 2), also shows a pronounced inflection at r~ = 2. This agrees with a more endothermic process for the formation of the third complex than for the second. From the data in Table 1, it follows that all the thorium acetate complexes are entropy stabilized, whereas the enthalpy change (positive) opposes their formation. This behaviour is in agreement with the nature of the acceptor and donor which are all "hard" in character[8, 9]. The values of AM for complexation reactions of hard acceptors with hard ligands generally decrease (they are less endothermic or even exothermic) as the energy of dehydration decreases for each consecutive step. Consequently, the entropy to be gained for each step should also decrease [1, |0].

2179

Thermodynamic properties of actinide complexes The marked increase in the value of AH~ at the third step of the thorium(IV)-acetate system is clearly connected with a change of coordination which can be also inferred from the increase in the entropy at this stage. Since such a change is likely to involve the expenditure of e\tra dehydration energy. Increases in the values of ,-XHiand ASj observed for the third step of the In'~-fluoride system[ll] and for the second step of the Zn-'~-acetate system [12, 13] have also been explained in terms of coordination changes. l.ike the thorium(lV)-acetate complexes, the uranyll VI)-acetate complexes (1) are stabilized predominantb by the entropy term. However, this term is not as great as for the thorium acetate complexes, probably on ~ccount of the difference in charge of the two metal ions. AckmJ~h,d,,,emep;t--'Fhis work has been supported by C.N.R.,

11a1\

REFERENCES

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