HYDROLYSIS OF BIS-(ACETYLACETONATO)-BERYLLICM(~~)
April, 1963
thallium salts demoiistrate iii a striking manner the special position of salts of non-octet) ions when they a c t as solvents in charge-unsymmetrical fused salt mixtures. Thus, we note from Fig. 3 that all the limiting heats in these solutions are much more positive than in the corresponding alkali nitrate systems. For all four systems the positive deviation is of the order of 2 kcal. In view of the uncertainty in the accepted ionic radii it may be fortuitous that the four values lie on or near a straight line which is roughly parallel to that for the alkali nitrate solutions. In our earlier discussion of the alkaline earth-alkali nitrate systems we noted that the small variations in the tern1 A in eq. 1 might be related to the polarieability of the divalent solute catioiis. However, the results obtained in the course of the present research seem to indicate that the principal contributions to the positive terms A , or more generally to the observed positive "shifts" in the limiting heats of solution, must be sought less in the properties of the solutes than in the properties of the soltent salts. The more positive enthalpies of solution in liquid silver and thallium nitrates compared to those observed in the corresponding alkali nitrates undoubtedly result from a combination
905
of several factors. Prominent among these presumably is the very significant van der Waals contribution to the cohesive energy in the solvent salts. It is worth noting that these positive contributions to the enthalpy of solution are significantly larger in the presently explored charge-unsymmetrical mixtures than in the simpler silver-alkali nitrate aiid thallium-alkali nitrate systems. For these solutions Blander6 was able to show that the magnitude of the observed positive shifts (compared to the binary alkali nitrates) is consistent with the change in the van der Waals interaction between second nearest neighbors. I t is possible that the larger positive shifts observed for solutions of divalent nitrates in the silver and thallium salts may be related to the local structural adjustment. This adjustment may, for example, involve both the population of ions in the various coordination shells and changes in interionic distances. In either case the van der Waals energy might be significantly changed. Acknowledgments -This work has been supported by The National Science Foundatioii (Grant 0 19513) and by the Office of Naval Research under Contract Number Nonr-2121 (11) with The University of Chicago.
HYDROLYSIS OF BIS-(ACETYLACETOKhT0)-BERYLLIUM(I1) BY R. W. GREENAND P. W. ALEXASDER School of Chemistry, S y d n e y (University, S y d n e y , Australia Received October 29, 1962 Distribution of 7Be between water and cyclohexane in the presence of excess acetylacetone a t 25' has been used to study the formation and hydrolysis of complexes between p H 4 and pH 11. At zero ionic strength, stability constants are log p1 7.96 and log PZ 14.67. The mono-complex, BeL(H20)*,undergoes two stages of acid dissociation with pki = 6.4 and pkz == 9.8.
Bis-(acetylacetona,to)-beryllium(II), first prepared by Combes, has attracted attention through its high stability and ready solubility in non-polar solvents. Step formation constants in water and in dioxane-water mixtures have been measured, and complete extraction of the bis complex into an organic from an aqueous phase at pH 5-10 has been r e p ~ r t e d . ~However, the last authors observed that the efficiency of extraction into chloroform fell sharply a t pH values greater than 10. The present paper describes the use of distribution measurements to investigate the nature and stability of the products formed' a t high pH. Experimental *s3
Cyclohexane, special grade for spectroscopy, supplied by British Drug Houses, was used without further purification. rlcetylacetone was purified by fractional distillation (b.p. 139"). Acidic aqueous solutions of acetylacetone were found to abeorb only weakly in the near-ultraviolet, but in alkaline solution a very strong absorption peak developed a t 294 mp. From measurements on a number of alkaline solutions, the molar extinction coefficient a t 25" was determined as 2.40 X 104, with a standard error of 0.01 X lo4. This value was then used for measurement (1) A. Combes, Compt. rend., 119, 1222 (1894). (2) R. M. Izatt, W. C. F'ernelius, and B. P. Block, J . Phys. Chem., 69, 80 (1955). (3) R. M. Izatt, W. C. Fernelius, C. G. Raas, and B. P. Block, ibzd.,69, 170 (1955). (4) J. A. Adam, E. Booth, J. I). 1%.Strickland, Anal. Chim. Acta, 6, 462
(1952).
of the acetylacetone concentration of experimental solutions.
Aqueous solutions were suitably diluted and made alkaline and their optical densities a t 294 mp were determined. Solutions in cyclohexane were completely extracted into aqueous alkali and treated similarly. I n neither instance did the presence of small amounts of beryllium mar the precision of the determination. Beryllium perchlorate solution was prepared from Analar beryllium sulfate and barium perchlorate. After filtration, the solution was evaporated t o crystallization on a steam-bath. Crystals of beryllium perchlorate were filtered and redissolved in water as an approximately 0.2 M solution. This was standardized by the chromazurol S method6 and then appropriately diluted for distribution experiments. Beryllium solutions were labeled with 7Be, supplied in carrier-free solution by the United Kingdom Btomic Energy Authority. The beryllium content of experimental solutions was related to known standards by counting the y r a y emission of 7Be with a thallium-activated sodium iodide well crystal coupled to a scintillation counter, Ekco Typo N664A. pH measurements were made a t 25' with a Radiometer 4 pH meter with saturated calomel and glass electrodes, standardized against 0.05 M potassium tetroxalate (pH 1.681), 0.05 M potassium hydrogen phthalate (pH 4.005), or 0.01 M sodium borate (pH 9.177).e For spectrophotometric measurements, a Hilger Uvispek instrument was used with I-cm. quarta cells thermostated a t 25". Distribution of solutes between water and cyclohexane was effected by gentle agitation in rotating glass cells immersed in a thermostat a t 25".
Results As a preliminary to examining the equilibria involving ( 5 ) L. 0 . Matveev a n d I. S. hIustafin, T i d y Komiss. Anal. Khzm. Akad. Nauk S S S R , 11, 217 (1960). ( 6 ) V. E. Bower a n d R. G . Bates, J . Res. Natl. Bur. Sld., 69, 261 (1957).
R. W. GREENAND P. W.ALEXANDER
906
Yol, 67
points, together with a distribution cuwe calculated from plc = 9.03 on the assumption that only the acid form (HL) is soluble in cyclohexane, with a distribution coefficient of 1.02. From 1 M NaC104 the distribution coefficient was 0.79. Distribution of beryllium between water and cyclohexane was established in less than 24 hr. at all pH values. Some of the experimental points are shown in Fig. 2, where the distribution coefficient, D = (total Be concn. in organic phase)/(total Be concn. in aqueous phase), is plotted against pH. I n these experiments total ligand concentrations in each phase were also measured. Figure 3 shows distribution of beryllium from 1.iV n’aC104 into cyclohexane in the presence of excess acetylacetone.
Discussion 3
7
5
9
PH.
Fig. 1.-Distribution of acetylacetone between water and cyclohexane. The curve ie calculated froin plz = 9.03 and d = 1.02,
4
d
2
The curves of Fig. 2 show that, in the presence of a sufficiently large excess of acetylacetone, D is almost constant in the pH range 5-10 and is independent of the total Be concentration. It is evident that polynuclear complexes are of no importance here. Since the system is clearly not saturated with beryllium, the constant distribution coefficient indicates a constant maximum value for the concentration of distributable species, which, considering the non-polar character of cyclohexane, can only be the bis complex, BeL2. I n this pH and concentration region, then, the beryllium must be virtuaIly all converted to this complex and D must now be identical with DZ = [BeLZ]orgsnia/ [ R ~ L z ] Its ~ ~ ~ ~ ~ ~ ~ value is 5.6, or 7.9 from 1 flfNaC104 solutions. From solutiom more acid than those represented by the plateau region of the curve, distribution is a function of acetylacetoiiate ion concentration IL-1, and hencc of pH. I n the absence of polynuclear complexes it is easily shown that
(D - Dz)PzY’ [L-12
0
4
8
12
PH.
Pig. 2.-Distribution of beryllium between water and cyclolo-’ 151; hexane in the presence of excess acetylacetone: 0,[LT], M ; [Berl, [&TI range, 6 x 10-6 to 3 x 10-2 X . e, [ L T ] , 6 X t o 6 X low3111. 8 , [LT], M; [&?TI,6 X 1O-’bl.
beryllium, the pk of acetylacetone a t 25” x a s determined. The conventional pH titration method with a 10-3 M solution, alter the application of an approximate activity correction of log y = -0.01,’ gave pk = 9.05with a standard error of 0.006. This was confirmed by a series of spectrophotometric determinations at 3.5 X 10-5 M , yielding pk 9.03 with a standard error of 0.008. I n 1 M h-aC104 solution the spectrophotometric method gal-e plc = 8.71. The distribution of acetylacetone between water and cyclohexane reached equilibrium in less than 5 hr., but equilibration was usually allowed to continue overnight. From acid solutions, the distribution coefficient, d = (concentration in cyclohexane) /(concentration in water), was found to be very nearly unity, but fell sharply near p H 8. Figure 1 shows experimental (7) E. Giintelberg, 2. physzk. Chem., 128, 1QQ (1926).
+ DPir [L-] + D
=
0
(1)
where pi, pz arc the therfhodynamic step stability constants and y is an approximation to the activity coefficient of a singly charged ion calculated from Guntelberg’s formula.’ Equation 1 is most easily applied to experimental points, like those of set A (Fig. 2), where the total ligand concentration, ILT], is a t least ten times as great as the total Be concentration, [BeTI, and the plateau region of the curve is fairly extensive. As a first approximation, the concentration of free ligand, [LF], can then be equated to [LT], and [L-I can be found from pH and pk, with appropriate correction for activity. Any of the usual curve-fitting techniques applied to eq. 1 then yields pi and ,&. From these and the above estimate of [L-1, estimates of [BeL] and [BeL2]can be formed and subtracted from [LT] to give a second approximation to [Lp]. This in turn leads to refined values of p1 and Pz. It was found in practice that one such repetition was sufficient for convergence to constant values of pi and P2. Table I reports two sets of results: those for solutions of ionic strength < 0.02, corrected for activity; and those for 1 M NaC104, which are stoichiometric constants. Agreement with the results of pH titration methods used by previous authors2 is satisfactory. Above pH 10 nearly all the free ligand is in the anionic form, insoluble in cyclohexane. Analysis of the cyclohexane phase in equilibrium with these alkaline
HYDROLYSIS OF
April, 1963
BIS-(ACETYLACETOKATO)-BERYLLIUM(~~)
TABLEI EQUILIBRIUM CONSTANTS Equilibrium
Dissnciation of HL Formation of BeL(H,0)2 Formation of BeLt Acid dissociation of BeL( Hz0)2
Constant
Ph log 81 log 82 Pkl PkL
Thermodynamic
9 03 7 96 14 67 6 4
9019
8 Stoichiometric (1 M NaC104)
8.71 7 55 14 35
6
9.s
solutions showed it to contain beryllium and acetylacetone in the exact proportion of 1:2, indicating that the only distributable species is still BeL2. The fall in D in this high pH region thus points to a fall in [BeL2]in the aqueous phase, suggesting the formation of hydroxy-complexes insoluble in the organic phase. Within the limits of the distribution technique, 1) is independent of [Beu], so that polynuclear complexes can be neglected, aind the most probable hydroxycomplexes can be formulated BeL(OH)H20 and BeL(OH)2. These two complexes appear as hydrolysis products of BeLz, but formally it is more convenient to treat them as products of the first and second acid dissociations of the mono complex, BeL(H,O),, with equilibrium coiistaiits
2
0
4
8
12
PH.
Fig. 3.-Distribution of beryllium between 1 M XaC104 and cyclohexane in the presence of excess acetylacetone: [LT],lo-' M ; [Be*], 6 X l O - 3 M . Broken curve reproduces results for water from curve A, Fig. 2.
If, as seems justified below p H 11 in the presence of excess ligand, we :now neglect water-soluble species containing no acetylacetone, it is possible to write for the aqueous phase [ B ~ T=] [BeL(HtO),+]
+ [BeL(OH)H*O]+
and an expression for the distribution coefficient
or
I n the presence of emess ligand, the left-hand side of eq. 2 can be estimated and plotted against 1/[H+] to give values for k, and kp, which can then be refined by successive approximlatioiis. The numerical values of plcl and pk, shown in, Table I were derived in this way. It should now be loossible to use the two formation constants, P1 and P p , and the two acid dissociation constants, kl and k,, together with d and D2,to calculate
D for any set of conditions. Successive approximations are again necessary but, with the fairly large excess of ligand used here, present no serious computational problem, and two cycles are sufficient. The results of these calculations are represented by the continuous curves of Fig. 2. The equilibrium constants were derived from the experimental points of set A, but calculation shows good agreement with observation within a hundredfold range of both [ B ~ Tand ] [LT]. Only for the most dilute solutions at low pH is there any noticeable discrepancy. This divergence does not arise from hydrolysis of Be2+, as can be shown by a more extensive calculation in which the hydrolysis constants of Kakihana and Sil16n8 are incorporated in the above treatment. It is possible that, under these extreme experimental conditions, our assumption of the constancy of Dz is not wholly justified. Otherwise the wide range of experimental results are satisfactorily accounted for by the interpretation adopted above. Acknowledgments.-The authors gratefully acknowledge a research grant from the Australian Atomic Energy Commission, and also the assistance of the Analytical Chemistry Section of the A.A.E.C., who standardized the beryllium solutions. (8) H. Kakihana and L. G. Sillen, Acta Chem. Seand., 10, 985 (1956).
