A spectrophotometric study of trivalent actinide complexes in solutions—III[1]

J inorg,nucl.Chem.,1969.Vol.3I. pp. 1807to 1814. PergamonPress. Printedin Great Britain

A SPECTROPHOTOMETRIC ACTINIDE COMPLEXES

STUDY OF TRIVALENT IN SOLUTIONS-Ill[I]

A M E R I C I U M WITH BROMIDE, I O D I D E , NITRATE AND CARBONATE LIGANDS[2] M. SHILOH*, M. GIVON and Y. M A R C U S t Israel Atomic Energy Commision Yavne, Israel

(Received I0 October 1968) A b s t r a c t - T h e light absorption spectra of americium(Ill) in concentrated aqueous lithium bromide, iodide and nitrate, magnesium iodide, potassium carbonate and nitric acid have been obtained, together with those of extracts into triisooctylammonium nitrate solutions in xylene. Effective stability constants for the first (inner sphere) complex are, log/3* = - 1.3 _+0.1 for AmN 032+ and --3.3--+ 0.1 for AmBr2+; the iodide complex is much weaker than the latter. Solubility studies show that [Am(OH)(CO3):j] 4- exists in 0-1-0-6 M potassium carbonate solutions. The spectral changes are discussed in terms of the environmental effects on the 5f-5ftransitions in americium(II I), INTRODUCTION

As A PART of an effort to elucidate the complex formation of the trivalent actinides in solution,[1,3,4] the chemistry of americium has been studied. The present paper deals with bromide, iodide, nitrate and carbonate ligands and the results for the chloride ligand will be reported in a subsequent publication[5]. Spectrophotometry has proved[l, 4] to be useful for the investigation of these systems since, in spite of being rather well shielded, the 5felectrons are affected by the environment, and do participate in bonding, even more than do the 4f electrons of the lanthanides [6, 7]. Cation exchange or solvent extraction studies indicate complex formation in americium(Ill) solutions containing 0-1 to 1 M chloride and nitrate, and even lower concentrations of sulfate. These, however, are solvent-shared ion pairs, since close contact of the metal ion and ligand, leading to coordinate bond formation and spectral changes, occurs only in much more concentrated solutions. In contrast to the trivalent oxidation state of the lighter actinides, americium(Ill) is quite stable. The high specific activity of the americium isotope available, 24'Am, however, limited the spectral region which could be studied, as well as the concentrations used. *Present address: Scientific Department, Israel Ministry of Defence, Tel-Aviv, Israel. tPlease send requests for reprints to this author, at his present address: Dept. of Inorganic and Analytical Chemistry, The Hebrew University, Jerusalem, Israel. 1. M. Shiloh and Y. Marcus, Neptunium and Plutonium- I I, J. inorg, nucl. Chem. 28, 2725 (1966). 2. Partly material from a Ph.D thesis submitted by M. Shiloh to the Hebrew University, Jerusalem (1964). Presented in part at the 149th A.C.S. Meeting, Detroit (1965). Preliminary publications in lsraelAEC Rep. IA-924 (1964); 1A-822, p. 67 (1962); 1A-1128, p. 99 (1966). 3. Y. Marcus, M. Givon and M. Shiloh, Proc. 3rd U.N. Int. Conf. Peaceful Uses Atom. Energy. Geneva, 1964, 10, 588 (1965). 4. M. Shiloh and Y. Marcus, lsraelJ. Chem. 3, 123 (1965) (Part 1). 5. Y. Marcus and M. Shiloh, (Part IV), lsraelJ. Chem. 7, in press (1969). 6. G. R. Choppin, D. E. Henrie and K. Buijs, lnorg. Chem. 5, 1743 (1966). 7. I. Abrahamer and Y. Marcus, lnorg. Chem. 6, 2103 (1967);J. inorg, nucl. Chem. 30, 1563 (196~). 1807

1808

M. S H I L O H , M. G I V O N and Y. M A R C U S EXPERIMENTAL

Americium obtained as 241AmO2 from the U.S. Atomic Energy Commission, was dissolved in the appropriate acids to form stock solutions of about 0-1 M Am(III). Concentrated solutions of lithium bromide, iodide and nitrate, magnesium and zinc iodides, potassium carbonate, and nitric acid were prepared from reagent grade chemicals. Triisooctylamine, from Carbide and Carbon Co. Inc. was dissolved in xylene, and equilibrated with 1 M nitric acid to form the ammonium nitrate salt (TIOAN). This was then preequilibrated with aqueous 8 M LiNOa. Absorption spectra were obtained with a Cary Model 14 spectrophotometer, with a cell compartment thermostated at 25°C. All aqueous solutions, except those containing nitrate, were placed over a layer of liquid zinc amalgam in the spectrophotometer cell, in order to avoid the formation of oxidation products, such as bromine, due to the intense a-activity. However, the absorption by nitrate is so intense that measurements could not be made below 320 nm in any case, so reduction of the solution was not required. Cells containing the most concentrated, hence viscous, solutions were centrifuged before the measurements in order to remove bubbles which caused "noise" in the spectra. The concentration of americium was determined by a-counting after dilution, using the half-life of 458 yr (specific activity of 1.73 x l0 a dpm//zmole).* All americium activities were handled in glove boxes, and a special box was used for a-containment of the cells inside the spectrophotometer. Lead shielding was used as required by the y-radiation of 241Am. The solubility of Am(III) in carbonate solutions was determined by introducing 25/xl of stock solution (pH = 3 in dilute perchloric acid) into 2.5 ml of potassium carbonate solution in a centrifuge cone. After standing 24 hr at room temperature with occasional shaking, the precipitate was centrifuged off and the americium in the supernate determined by counting. The ratio of carbonate to americium in the precipitate was determined by liberating COz from it by excess 1 : 1 H2504, leading it with an argon stream (to avoid contamination by CO2 from the air) into excess factorized Ba(OH)2 solution, and back titrating the latter. The americium content, ca. 3-5/~moles in each experiment, was determined spectrophotometrically in the residual sulfuric acid solution. RESULTS

Hydrated americium(Ill) ions, i.e. solutions in perchloric acid or other dilute acids, have two prominent bands in the visible region. One of these is at 19880 cm -1, henceforth called the 503 nm band, with e = 380 M - l c m -1, and the other is at 12320 cm -1, henceforth called the 812 nm band, with e = 60 M - l c m -1. Both bands are sensitive to the environment, being shifted or split by the presence of suitable ligands. Bromide solutions. Like the lighter actinides [ 1,4], americium(I I I) has a characteristic 5 f ~ ~ 5 f ~-1 6d transition of high intensity in concentrated halide solutions. In chloride solutions this occurs at 4 2 5 0 0 cm -~ (235 nm)[3, 5], and is expected to shift to 4 2 0 0 0 cm -1 (238 nm) in bromide solutions. However, the absorption edge of the bromide ion itself has ~ = 1.4 M - l c m -1 at this wavelength in lithium bromide[8], and since this salt must be present at very high concentrations, it precludes the observation of this band. In the visible region the 503 nm band is sensitive to the environment and, in concentrated bromide solutions shifts to 19780 cm -~ (505.5 nm), and decreases in intensity, to E = 260 M -1 cm -1, in the most concentrated solution studied, 11.44 M LiBr (Fig. 1). The 812 nm band is also shifted to lower energies, 12200 cm -~ (820 nm). * A recent redetermination of the half-life of 241Am has yielded a value of 433 yr. (see K. W. Bagnall et al. J. chem. Soc. (A), 133, 1968). This results in a 6% increase in specific activity and based on this

value the solubilities measured by c~-counting would be 6% lower and the molar absorptivities 6% higher than those reported here. 8. Thanks are due to Mrs. S. Skurnik, who provided data on the absorbance in concentrated alkali bromide solutions.

Trivalent actinide complexes in solutions

500

20000

19800

[

i

19600

1809 19400

i

cm-

I

C

400

300--

~~

200

~ 11'

LI'

,oo 500

/,,,_.~

"',,

".....

505

510

515

nm

520

Fig. 1. The 503 nm band in the americium(Ill) s p e c t r u m : - - - - I M HCIO4, - - - 11.4 M LiBr, - - - - 6.0 M K2CO3 solutions. The molar absorbance E in M 1 cm 1 of 0.6-2-0 mM americium solutions is plotted against the wavelength.

The molar absorptivity at 503 nm, A, could be used to calculate the complex formation in 8.74-11.44 M LiBr solutions. The effective activity of the bromide [4, 9], a = CLi~rYLmr,changes in this region from 330 to 2580, i.e. about eightfold. Extrapolation of A to 1/a = 0 gave Ao~= 200 M - ' c m -1, and from the known A0, the function (A o - A )/(A --A ~) was calculated, and found to be linear with a. The slope thus equals#* = ( 5 . 3 _ 1.4) x 10 -4 M -1, signifying formation of the species A m B r z+. Changes in the absorptivity for the other bands were too slight to give a value for the stability constant. Iodide solutions. The absorption spectrum of americium(Ill) in saturated lithium iodide (7-37 M) or zinc iodide (5.6 M, or 11.2 N in iodide) is the same as that in dilute aqueous solutions. Only in concentrated magnesium iodide solutions (4.1 M, or 8 . 2 N in iodide, where a=clY+_Mgi2 ~ 2100) is the water activity sufficiently low, and the effective iodide activity sufficiently high for spectral changes to be observable. As in bromide solutions, there is a shift of the 503 nm band towards i 9760 cm -1 ( 5 0 6 rim). Nitrate solutions. The absorption spectrum of americium(III) in nitrate solutions is shown in Fig. 2. In aqueous 8 M lithium nitrate or nitric acid, there is a decrease in intensity of the 503 nm band, with a distinct shoulder appearing in lithium nitrate at 19400 cm -1 (515 nm). The 812 nm band, however, changes its shape and shifts towards higher energies, contrary to the behaviour in halide solutions. Although nitric acid is only partly dissociated, the mean ionic activity coefficients are higher in nitric acid solutions [ 10] than in lithium nitrate solutions [ 111, 9. Y. Marcus, Record Chem. Progr. 27, 112 (1966). 10. W. Davis, Jr. and H. J. DeBruin, J. inorg, nucl. Chem. 26, 1069 (1964). 11. M. Gazith, IsraelAEC Rep. IA-1004 (1964).

1810

M. SHILOH, M. G I V O N and Y. M A R C U S 24000

400

I

22000 I

20000

18000

I,

13000

I ~#¢

12000 1

= =

6

!1": il

,oo-

~1~ il

"

il 0 400

cm -I

I

450

500

'

550

750

800

,,'... k nm

850

Fig. 2. The 503 and 812 nm bands in the americium(Ill) spectrum: . . . . . . 1 M HC104, - - - 8.0 M LiNO3, 20% v/v triisooctylammonium nitrate in xylene. The molar absorbance ~ in M -a cm -~ of 0-5-1 '0 mM americium solutions is plotted against the wavelength.

and at equi-molar concentrations the effective nitrate activity a = CNO3Y_+(no~ Li)N03 is comparable in the two solutions. Thus, at 8 M, a = 28.5 in lithium nitrate solutions, and 22.3 in nitric acid. H o w e v e r , the solubility of the latter is considerably higher, so that an effective activity a = 114 is attainable at 1 5 M nitric acid, much b e y o n d the range for lithium nitrate. Results were obtained for both electrolytes, but because of the wider range, those for nitric acid are better. An isosbestic point (Fig. 3) is obtained at 798 nm, and is taken as evidence that only two absorbing species are at equilibrium, with due cognizance of the limitations of this argument [ 12]. T h e spectrum could be extrapolated to infinite nitrate concentration (1/a = 0), as shown in Fig. 3. Analysis of the data, as above for bromide, showed that (Ao--A)/(A--Aoo) =/3*a, with /3" = ( 5 . 0 + 1-0) × 10-2 M - ' , signifying the formation of A m N O 2+. T h e spectrum observed for an extract of americium(III) into 20% v/v triisooctylammonium nitrate in xylene shows shifts even more pronounced than those in aqueous solutions. T h e 503 nm band is split into at least three bands, at 19800, 19550, and 19230 cm -1, the middle one being the most intense, while the 812 nm 12. J. Brynstead and G. P. Smith,J. phys. Chem. 72,296 (1968).

T r i v a l e n t a c t i n i d e c o m p l e x e s in s o l u t i o n s I

I

I

I

181 I

!

[

sol

e o

40 20 740

.""'""" "'--.... ,,

,j~//

"-..',

780

820

nm

860

Fig. 3. T h e 812 n m b a n d of a m e r i c i u m ( I l l ) in nitrate s o l u t i o n s : 1 M HCIO4, - - 7.7 H N O 3 , - - - - 15.0 M H N O 3 . . . . . . . . . . e x t r a p o l a t e d to infinite e f f e c t i v e nitrate a c t i v i t y . T h e m o l a r a b s o r b a n c e • in M - ' cm-a is p l o t t e d a g a i n s t the w a v e l e n g t h .

band shows two bands at 12740 and 12500 cm 1 (Fig. 2). The small peak at 21900 cm -1 is strongly enhanced in this solutions, attaining • = 20, which is much higher than the values for other americium(Ill) bands in this region. Carbonate solutions. In potassium carbonate solutions the complexation is stronger than in halide or nitrate solutions, as shown by the shift of the 503 nm band to 19660 c m - ' in 6 M potassium carbonate. The gradual shift of the band at increasing concentrations indicates the formation of several species. In the range 0-12-0-60 M potassium carbonate, however, only one species predominates, as is seen from the constancy of the molar absorbancy of the peak at 19700 cm-' and from the constant slope of two in the double logarithmic plot of the americium solubility SAm and the potassium carbonate concentration (CO3"-) (Fig. 4). In several experiments, the ratio of carbonate to americium in the precipitate was found to be 1-60, so the composition Am2(CO3)3 was assumed. To explain the slope of two, it is necessary to assume the formation of the species [Am(OH)(CO.~)3p- in the concentration range studied, assuming the following equilibria Am,,(CO~)a(s) ~ 2Am a+ + 3CO32-

K,

(1)

Am :~++ O H - + 3CO:~2- ~-- [Am(OH)(CO:03] 4-

Ks

(2)

CO:('- + 2H20 ~.~ 2 O H - + H,,CO3

Ks

(3)

Equilibrium (3) was found to show constant activity for HzCO3 in the range studied (Fig. 4), so that the equilibria can be summarized by SAm = ([Am(OH)(CO3)3] 4-) = K,1/2K2K3'/2(H.,COa)-'/2(CO:~2-) 2 = K'(CO.~2-) 2 (4) as found experimentally, with K' = 1.8 × 10-aM -'. In this calculation concentrations were used instead of activities. This is a crude approximation particularly for equilibria involving highly charged ions.

1812

M. S H I L O H , M. G I V O N and Y. M A R C U S I -

[

I

3.0

pH 12.5

E ~o

12.0

o -4

11.5 f

O

O

-4.5

o

-i.o .

-05 ~ log c

IL° KzCO 3

Fig. 4. The solubility of americium(l I I) in (left-hand ordinate), and the pH of (right-hand ordinate), potassium carbonate solutions. DISCUSSION

The stabilities of the trivalent chloride and bromide species, MX 2+, of uranium, neptunium, plutonium and americium have been compared previously[l]. A nearly linear positive dependence of log fll* on the atomic number is noted, except for neptunium, which shows some excess stability. For iodide, the data for neptunium and plutonium in concentrated magnesium iodide solutions show no spectral changes [1], while for americium there is some indication of complexation. Thus with iodide the order of complexation is also Np, Pu < Am. It is not possible to compare the nitrate complexes in the same way, since only those for plutonium and americium would be stable to oxidation. For the former, the qualitative observation of more pronounced spectral effects than in halide solutions, hence presumably also stronger complexing, has been made. For the latter this has been confirmed quantitatively, the strength of the complexes being NO3- > C1- > Br- > I-. Care must be exercised in comparing effective stability constants for different ligands, particularly if they were obtained from data for different spectral transitions. The nitrate ligand, in particular, can be misleading [6], and the question of inner-sphere vs. outer-sphere complexing [6, 7, 13] should be carefully considered. 13. J. C. Barnes, J. chem. Soc. 3880 (1964).

Trivalent actinide complexes in solutions

1813

In the present case the stability constant for chloride[2, 5] were obtained from two dissimilar spectral transitions: 5f 6 -~ 5f 5 6dfor the 235 nm band and 5f 6 -~ 5f 6 for the 503 nm band, with good internal agreement. For the nitrate, the constant obtained from the results for the 812 nm band also describes well the data for the 503 nm band, although here there is no shift in peak position, only a decrease in intensity and the growth of a shoulder. Since for the complexes discussed here f - f transitions are affected, it is concluded that they are inner type complexes. It is of interest to compare their stability with those of outer-sphere complexes (solvent-separated ion pairs) studied by extraction and ion exchange methods. It is found that for those cases where comparable data exist (PuC12+, PuCI~+, AmCI 2+, AmCl2 +, AmNO:} +) the outer-sphere stability constant is about 100 times higher, per ligand, than the innersphere constant. In other words, there is an increase of ca. 2-7 kcal in the free energy as a water molecule is eliminated when a ligand moves from an outersphere to an inner-sphere position. Both the enthalpy and the entropy changes of complexation are negative for outer-sphere and positive for inner-sphere complexes[14] but it is not possible to obtain from these data the values for the replacement of a water molecule, since they pertain to different ligands. Both the 503 and the 812 nm bands belong to transitions to degenerate energy levels of high multiplicity. Lowering of the symmetry around the americium ion will cause partial removal of the degeneracy and of the restriction on inner orbital transitions. Thus, in addition to shifts in the bands, due to ligand field effects, there may also be changes in the intensity of the several components of a multiple band, which will result in changes of shape, shift of the maximum, appearance of shoulders, or in extreme cases, the splitting of the band. In this respect, the bromide, iodide and carbonate ligands affect the spectrum differently from the nitrate and the thiocyanate[15]. The ground level of americium(III) has been assigned as a 7F0term[16] and the 812 iam band is due mainly to a 7F0--~ 7F6 transition, with some mixture from the first excited level with J = 0. This band is shifted to lower energies in bromide, iodide and carbonate solutions, but in nitrate, the J = 0 component, which lies at higher energies than the ground J = 6 state [16], is enhanced, resulting in a net shift to higher energies, as already noted by Carnall et al. [ 18]. The 503 nm band, due to a transition to the first excited level with J = 6 (assigned as ~G~ by Gorbenko-Germanov[16] but later as '~L6 by Carnall and Wybourne[17]) is much more strongly affected by the ligands, in particular as regards the intensity. The splitting of the band (Fig. 2) observed for the extracts into the long-chain ammonium nitrate solutions shows that here the americium is in an environment very different from that in aqueous solutions. The changes, however, are not as pronounced as those observed for extracts into long-chain 14. G. R. Choppin and W. F. Strazik, lnorg. Chem. 4, 1250 (1965); G. R. Choppin and J. Ketels, J. inorg, nucl. Chem. 27, 1335 (1965); G. R. de Carvalho and G. R. Choppin, ibid. 29,725 (1967). 15. Y. Marcus and E. Yanir, Unpublished results, 1968. 16. D. S. Gorbenko-Germanov, Fiz. Probl. Spektr. Akad. Nauk SSSR. Kom. Spektr. Moscow. Materialy 13-.go Sovescbaniya, Leningrad (1960). Vol. 1, pp. 242-247, (1962). 17. W. T. Carnall and B. G. Wybourne,J. chem. Phys. 40, 3428 0964). 18. W. T. Carnall, P. R. Fields and B. G. Wybourne, J. chem. Phys. 41, 2195 (1964). 19. Y. Marcus, IsraelAEC Rep. IA-1020 (1965); IA-1021, p. 55 (1965).

1814

M. SHILOH, M. GIVON and Y. MARCUS

ammonium chloride solmions [5, 19] and certain solid chloride salts [5, 20], as well as for solid sulfates and double sulfates[21]. It should also be noted that the peak at 21900 cm -1, which is strongly enhanced in all these cases relative to the 503 nm band (and also absolutely), is "hypersensitive", involving a transition with A J = 2118]. Again, the process of relative enhancement is not as advanced in the nitrate as in the other cases quoted. T h e microsymmetry around the metal ions in long-chain ammonium nitrate extracts has been discussed for the lanthanides [7], and should be similar for americium, which shows the same second-power extractant concentration dependence as the lanthanides at tracer concentrations [22]. A coordination n u m b e r of six, one water molecule and five monodentate nitrate ions, is consistent with these spectral and extraction data, but other interpretations are possible [7]. T h e species proposed for americium(III) in carbonate solutions require comment. We find [Am(OH)(CO3)3] 4- in the range 0.1-0.6 M K2COa. F o r lanthanides, Sherry and Marinsky[23] report, from anion exchange data, that the species M(CO3)2-, M(CO~)33+ and M(OH)(CO3)4 ~- are formed successively, as the potassium carbonate concentration increases from very low values to 3 M. T h e species now proposed for americium(Ill) thus fits well into this series, being present in the intermediary concentration range, and is consistent with anodic electromigration data [24]. Acknowledgements - W e thank Mr. L. Cohen and Miss N. Abel for technical assistance.

20. Private communicationfrom J. L. Ryan, June 15 (1966). 21. G. N. Yakovlev, D. S. Gorbenko-Germanov, R. A. Zenkova, V. M. Razbitnoiand K. S. Kazanskii, Zh. Obshch. Khim. 28, 2624 (1958). 22. Y. Marcus, M. Givon and G. R. Choppin, J. inorg, nucl. Chem. 25, 1457 (1963). 23. H. S. Sherry andJ. A. Marinsky, lnorg. Chem. 3,330 (1964). 24. I.A. Lebedev, S. V. Pirozhkov and G. N. Yakovlev, Radiokhimiya 2,549 (i 960).