Detection of metal ion hydrolysis by coagulation. VI. Beryllium

JOURNAL OF COLLOID SCIENCE ~0, 322-329 (1965)

D E T E C T I O N OF METAL ION H Y D R O L Y S I S BY COAGULATION VI. 1 BERYLLIUM 2 Egon Matijevid Department of Chemistry, Clarkson College, Potsdam, New York Received September 30, 1964 ABSTRACT The hydrolysis of beryllium ion was i n v e s t i g a t e d b y means of the coagulation method. T h e charge of t h e the hydrolyzed species is found to be ~ 3 . This finding has been compared with the results of Silldn and collaborators and the differences have been discussed. INTRODUCTION

Hydrolysis effects of beryllium salts were first recorded over a century ago (1-4). Some of the work that followed contained more quantitative measurements, but little information regarding the composition of the hydrolyzed beryllium ionic species (5-9). A number of investigators, using various experimental techniques, have attempted to resolve the latter problem. Unlike the case of some other metal ions, the investigations on the hydrolysis of beryllium ion have led to consistent conclusions with regard to the following major points: (a) the ratio Be:OH in the hydrolyzed species is 1 : 1 (5, 10-21), (b) the complex species is polynuclear (10-32), and (c) no precipitate of beryllium hydroxide is formed unless the amount of added base is well in excess of 1 OH- per Be+e (6, 17, 19, 21, 27, 32). It follows from these experiments that soluble beryllium hydrolyzed species must correspond to the general formula of Be.(OH). ~+. The only exception to this formulation is found in Mattock's work (29), where it is concluded that the hydrolysis of beryllium involves the formation of a large number of polynuclear species, and ill the work of Sidgwick and Lewis (23), in which they suggested the formation of a pentameric species of the type [Be.4BeO] ~+. If the general formulation Be~(OH)~"+ is accepted, the species becomes defined if n is known. It is also generally recognized that a single hydrolyzed beryllium species is predominant. A number of investigators (11, 14, 24, 30, 32, 33) favored the valence n = 2, thus proposing dimerized beryllium ' P a r t V, E. Matijevi6, K. G. M a t h a i , and M. Kerker, J. Phys. Chem. 66, 1799 (1962). 2 Supported by a g r a n t from the A r m y Research Office No. ARO-(D)-31-124-G320. 322

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ions Be2(OH)22+ or Be202+. Others (13,17,21,25-27) claimed, on the other hand, tetrameric species Be4(OH)44+. Finally, the very careful work of the Swedish school (18-20,34,35) seemed to give results compatible with the assumption of a trimeric formulation Be3(OH)33+ as the predominant species with a small amount of Be2OH3+. At low total beryllium salt concentration and at higher pH Be(OH)2 was also indicated. Garrett then recalculated his own solubility data (30) and found that these can be explained better if Be~(OH)~3+ is assumed as the hydrolyzed species (31). It appears that an independent determination would be desirable to resolve the problem of the composition of the hydrolyzed beryllium ion. In view of the general acceptance of the formulation Be~(OH)~n+, a determination of the actual charge of the hydrolyzed ion would give the decisive answer. We have shown earlier that the coagulation method can provide this information in a very elegant way (36-39). In this work we have determined the charges of the hydrolyzed beryllium ion and shown that it is +3. It also appears from the critical coagulation concentration of the hydrolyzed ion that the predominant species is trimeric. EXPERIMENTAL

We have followed the coagulation of silver bromide in statu nascendi by mixing equal volumes (5 ml.) of silver nitrate solutions of constant concentration with solutions containing constant amounts of potassium bromide component aRd varying amounts of beryllium nitrate. The turbidity changes were measured by observing the light scattering at 45 ° in an Aminco light-scattering microphotometer using the 546 mt~ line. Nitric acid, when required, was added to the solution of bromide and beryllium components. However, when measurements were performed in the presence of NaOH, the silver nitrate and beryllium nitrate solutions were mixed first, and the NaOH was added to the bromide solutions. This was done ia order to prevent the formation of silver or beryllium hydroxide owing to the local supersaturations when the more concentrated stock solutions of NaOH were added. All measurements were performed at 25°C. For a detailed description of the experimental technique and the method of analyzing the data iu order to obtain the critical coagulation concentration we refer to our previous papers (36-38, 40, 41). Half of the total amount of the prepared sol solution was used for turbidity measurements whereas the rest was utilized for pH measurements. A Beckman Model G pH meter was employed, the scale of which was calibrated with appropriate buffer solutions prior to taking pH readings. All chemicals were of the highest purity grades further purified when necessary. Beryllium perchlorate was prepared and analyzed as described by Kakihana and Sill~n (18). The concentration of beryllium in the stock solutions was determined gravimetrically. Water was doubly distilled, the

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MATIJEVIC

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FIG. 1. Plot of the critical coagulation concentration (c.c.c.) against pH for beryllium nitrate (O) and beryllium perchlorate (D) for a silver bromide sol i n s t at u n a s c e n d i . Concentrations: AgNOa, 2 X 10-4 M; KBr, 2 X 10-3 M; HNO3 or NaOH, varied. Full circles correspond to the 3[Be3(OH)3a+]/[Be(NO~)2]tot~l vs. pH plot using log K33 = -8.66 (Kakihana and Sill~n) (18). Dashed line gives corresponding "log c.c.c. - p H " plot for zinc nitrate (36). For the latter plot the ordinate represents log molar conc. of Zn(NOa)~.

second time from an all-Pyrex still, and regularly tested for pH and conductivity. All glassware was thoroughly cleaned and steamed before use. RESULTS

The dependence of the critical coagulation concentration (c.c.c.) of beryllium ions on the pH of the medium is given in Fig. 1. A negative silver bromide sol in statu nascendi was used in these experiments. All concentrations are given in the legend. Circled points were obtained using beryllium nitrate as the coagulating electrolyte. To check whether there is any effect of nitrate ion upon the coagulation a few experiments were carried out with beryllium perchlorate, since perchlorate ion is usually considered to show least tendency to form complexes. These experiments are denoted with squares, and it is obvious that excellent agreement is found

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with data obtained using beryllium nitrate. Prytz (42) showed earlier that bromide ions do not form complexes with beryllium. Thus coagulation effects can be interpreted in terms of hydroxylation of beryllium ion at higher pH's. In order to ascertain whether insoluble beryllium hydroxide is formed in any of the investigated systems, blank experiments were carried out which contained all components with the exception of silver nitrate. No precipitation of beryllium hydroxide was observed in any of the experiments recorded in Fig. 1. Aging of beryllium nitrate solutions did not influence the c.c.c. The dashed line in Fig. 1 indicates the c.c.c, vs. pH plot for zinc iou taken from earlier work (36). This will be utilized later on in the discussion. Discussion

The coagulation concentration of beryllium ion remains constant up to pH ~-~ 3. It decreases then sharply and levels off above pH 5.5. According to the Schulze-Hardy rule a decrease in coagulation concentration signifies an increase in the charge of the coagulating ion. Earlier, we had determined experimentally the relationship between the c.e.c, and the charge of the counter]on for this particular sol (43-45). From this it is easily deduced that the hydrolyzed beryllium ion carries a -{-3 charge. Actually the c.c.c. of the hydrolyzed beryllium ion is somewhat lower than the value reported for ions of charge + 3 (45). However, the coagulation concentrations tabulated in the paper cited, refer to ions when their size is extrapolated to zero. It is known that the c.c.c, decreases with increasing counter ion size for ions of the same charge. A beryllium hydrolyzed complex of charge 4 3 would be an ion of considerable size and this is why the c.c.c, is expected to be lower than for a simple trivalent ion. Sill~n and collaborators have concluded that n = 3 in Be~(OH)~ ~+, giving the predominant species Be3(OH)~~+ (18, 19, 34). These authors also found evidence for small amounts of Be2(OH) 3+. In addition, at low beryllium salt concentrations ( < 1 mM) and at higher pH (>5) they assumed the existence of the uncharged Be(OH)2 species. The following equilibria were postulated (18, 35): 3Be~+ + 3H20 ~ Be~(OH)38+ + 3H + 2Be~+ + H20 ~- Bc2(OH) 3+ + H + Be2+ + 2H20 ~ Be(OH)2 + 2H +

log K33 = -8.66 log K12 = -3.22 log K21 = -10.87

From these values for K33 and K21 (and with the second equilibrium constant neglected), concentrations of Be3(OH)33+, Be(OH)2, and Be2+ species were calculated for the conditions used in the coagulation experiments and given in Table I. It appears that at concentrations of Be(NO3)~ < 6 X 10+4//and at pH > 5.5 Be(OH)2 becomes the predominant species.

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MATIJEVI6 TABLE I

Concentrations of Be(OH)~+, Be(OH)s, and Be++ Species in Moles~Liter in Solution at Conditions Used in Coagulation Experiments Assuming Equilibrium Constants K33 and K21 (Columns ~, 4, and 5) and Assuming Only K33 (Column 6)

[Be(NO3)2] 1.6 X 1.6 X 1.25X 6 X 1.4 X 3 X 3 X 3 X

10 -8 10 -3 10 -8 10 -4 10 -4 10 -5 10 - s 10 -5

pH

[Be3(OH)]+]

2.0 3.0 4.0 5.0 5.5 6.0 7.0 8.0

9 9 4.1 7.7 8.8 2.3 2.9 3

X X X X X X X X

10 -is 10 -9 10 -6 10-~ 10 -6 10 -8 10 -16 10 -14

[Be2+] 1.6 1.6 1.24 3.3 5.0 2.2 2.4 2.4

X X X X × X X ×

10-* 10 -8 10 -8 10 -4 10 -6 10 -6 10 -6 10 -16

[Be(OH)~] [Be3(OH)] +] 2 2 1.6 4.1 6.3 2.8 3 3

X X X X X X X X

10 -1° 10 -8 10 -6 10 -5 10 -6 10 -5 10 -6 10 -5

9 9 4.15 8.7 2.3 5.5 9.5 9.9

X X X X X X X X

10 -12 10 -9 10 -6 10-~ 10 -5 10 -6 10 -8 10 -6

This is inconsistent with the coagulation experiments which indicate an increase rather than a decrease in ionic charge for the hydrolyzed beryllium species. The formation of lower charged, or even uncharged, species would lead to a sharp increase in the c.c.c., as was observed with some other ions (37, 39, 46). It is interesting to note that if only the formation of Be3(OH)38+ is considered and the concentrations are calculated using K33 (Table I, column 6) the agreement between Sill~n's data and the coagulation results becomes rather good. This is shown in Fig. 1, where 3[Be3(OH)33+]/[Be(NO3)~]Tor vs. pH (full circles) is plotted. The presence of small amounts of Be~(OH) 3+ could slightly influence this comparison. Since the charge is the same the coagulation method does not permit the detection of the latter species. The discrepancy between the coagulation data and the results of Sill~n and collaborators is difficult to resolve. The upper concentration of beryl]ium salt in coagulation experiments is just about equal to the lowest concentration used by the Swedish school. The question of the existence of some other complexes at lower concentration is still unresolved. However, it is unlikely that these complexes would be polymeric species of high charges which could compete with the lower charged or uncharged species. On the other hand, one could assume that Be(OH)~ may be strongly adsorbed on the negative silver bromide particles and thus induce a coagulation by adsorption. It is difficult to visualize why an uncharged hydroxylated species would be more strongly adsorbed than a charged hydrolyzed species. Experimental evidence obtained for a number of metal counter ions indicates that the Schulze-Hardy rule is valid when these ions are hydrolyzed. How much of these ions are adsorbed below and above c.c.c, as a function of pH is the subject of our present studies. It is hoped

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ION HYDROLYSIS

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that these experiments will illuminate the problems associated with the adsorption of counter ions more fully. The results with beryllium are in agreement with our previous findings on the hydrolysis of zinc ions (36). We found that the charge of hydrolyzed zinc ions is also + 3 and therefore concluded that the species must be polymerized. Although Perrin (47), on the basis of his recent potentiometric work, still assumes the existence of the simple ZnOH + ion, the refined method as employed by Biedermann (48) coidirmed our findings that hydrolyzed zinc ion is polymerized in the form of Zn2OHa+. In Fig. 1 we have given for comparison purposes the "c.c.c.-pH" curve of zinc nitrate. It is obvious that in the concentration range of our experiments zinc is less strongly hydrolyzed than beryllium. Also the c.c.c, of hydrolyzed zinc ion seems to be somewhat lower than the corresponding value for Be. If the total metal ion concentration plotted on the ordinate is divided by 2 for zinc and by 3 for beryllium approximately the same final value of c.c.c, for both ions is obtained. Since hydrolyzed zinc ion is a dimer, it would appear that hydrolyzed beryllium ion must be atrimer. Coagulation data indicate that the formulation of the hydrolyzed beryllium ion in the range of concentrations and pH used in these coagulation experiments is Bea(OH)a3+. The same predominant beryllium hydrolyzed species are assumed by Sill~n (19) but for higher concentration range and lower pH. It is not possible to explain this discrepancy at present. Our findings that beryllium nitrate and beryllium perchlorate solutions give the same results are significant from still another point of view. The potentiometric method requires addition of a high concentration of sodium perchlorate and as pointed out in the case of thorium ion (49) " . . . . i t would not be possible to distinguish between formulas with varying amounts of solvent and medium ions: X'r~h [Th2(OH)2 ~+] = ~ ~_~z_~_,t'i'

~" C1, (4q-2-~),, ~ H -2("H2u)~ j .

Thus, the potentiometric method does not give the charge directly, but this is computed assuming that no medium ions are included in the complex. On the other hand, the coagulation method, under favorable conditions, directly gives the charge. When this charge agrees with the computed value as obtained from other experimental methods, it is fair to conclude that no complexing has taken place with ions other than hydroxide. For the potentiometric work, the concentration of perchlorate ions is rather large. Unfortunately, the coagulation method cannot be used at such ionic strengths. Thus, on the basis of charge determination, a direct test whether at these high concentrations the perchlorate ion is incorporated in the hydrolyzed complex or not is not experimentally feasible. Finally, we have also been able to show that the hydrolyzed beryllium ions reverse the charge of the lyophobic colloid in agreement with our

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earlier observations (50). A detailed study of the reversal of charge by beryllium ion is forthcoming. It is worth while noting here that the filtration studies of Schweitzer and Nehls (28) can be explained on the basis of the reversal of charge effects. It would appear that the removal of beryllium ions by filtration is not due to the formation of colloidal beryllium hydroxide but rather to the adsorption of hydrolyzed beryllium species on the filter. The pH at which a sharp increase in the amount of beryllium removed by filtration is observed, corresponds to the region at which hydrolyzed beryllium reverses the charge due to adsorption on oppositely charged surfaces. ACKNOWLEDGMENT The assistance of Messrs. Christopher R. Davies and Ronald L. Nemeth who performed the experiments is gratefully acknowledged. REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28.

29. 30.

BERZELIUS, J. J., Schweiggers J. 15, 296 (1815). LE¥, H., Z. Physik. Chem. 30, 193 (1899). BRUNER, L., Z. Physik. Chem. 32, 133 (1900). PARSONS, C. L., J. Am. Chem. Soc. 26, 1433 (1904). BRITTON, H. W. S., J. Chem. Soc. 127, 2120 (1925). (~uPR, V., Collection Czech. Chem. Commun. 1, 377 (1929). BROWN, H. F., AND CRANSTON, J. A., J. Chem. Soc. 1940, 578. PRYTZ, M., Trans. Faraday Soc. 24, 281 (1928). LINCH, L. S., AND MOGILEY,M. E., Zhur. Obshchei Khim. 26, 331 (1956), English Translation. PRYTZ, M., Z. Anorg. Allgem. Chem. 180, 355 (1929). PUC~E, F., AND JOSIEN, M. L., Bull. Soc. Chim. France 1940, 755. JOSIEN, M. :L., Bull. Soc. Chim. France 1940, 955. SCHAAL,R., AND FAUCHERRE, J., Bull. Soc. Chim. France 1947, 927. HALDAR, B. C., J. Indian Chem. Soc. 25, 439 (1948). HEUKESHOVEN, W., AND WlN~EL, A., Z. Anorg. Allgcm. Chem. 213, 1 (1933). FELDMAN, I., AND HAVILL, J. R., J. Am. Chem. Soc. 74, 2337 (1952). FAVCHERRE,J., Bull. Soc. Chim. France 1953, 1117; ibid. 1954, 253. KAKIHANA,H., AND SILL]~N, L. G., Acta Chem. Scand. 10, 985 (1956). SILL~N, L. G., Quart. Revs. (London) 13, 146 (1959). CARELL, B., AND OLIN, /~k.,Acta Chem. Seand. 16, 1875 (1961). JANDER, G., AND J ~ R , K. F., KoUoid-Beih. 43, 295 (1936). PARSONS, C. L., ROBINSON, W. O., AND FULLER, C. T., J. Phys. Chem. 11, 651 (1907). SIDGWICK,N. V., AND LEWIS, N. B., J. Chem. Soc. 1926, 1287. PRYTZ, M., Z. Anorg. Allgem. Chem. 197, 103 (1931). FAUCHERRE, J., AND SCHAAL,R., Compt. Rend. 225, 118 (1947). SOUCHAY, P., Bull. Soc. Chim. France 1948, 143. TEYSSI~IDRE,M., AND SOUCHAY,P., Bull. Soc. Chim. France 1951, 945. SCHWEITZER,G. g . , ANDNEHLS, T. W., J. Am. Chem. Soc. 75, 4354 (1953). MATTOCK, G., J. Am. Chem. Soc. 76, 4835 (1954). GILBERT, R. A., AND GARRETT, A. B., J. Am. Chem. Soc. 78, 5501 (1956).

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31. SCHINDLER,P., AND GARRETT, A. B., Helv. Chim. Acta 43, 2176 (1960). 32. BJERRU~I, J., "Metal Ammine Formation in Aqueous Solution," p. 173 ft. P. Haase and Son, Copenhagen, 1957. 33. HONDA, M., J. Chem. Soc. Japan Pure Chem. Sect. 72, 361 (1951). 34. CARELL, B., AND OLIN, /~k., Acta Chem. Scand. 16, 2357 (1962). 35. HII~TANEN,S., AND SILL~N, L. G., Acta Chem. Seand. 18, 843 (1964). 36. 1V[ATIJEVI~, E., COUCH, J. P., AND KERKER, M., J. Phys. Chem. 66, 111 (1962). 37. MATIJEVI¢, E., MXTHAI, K. G., OTTEWILL, ]:~. i . , AND KERKER, M., J. Phys. Chem. 65, 826 (1961). 38. MATIJEVI(~, E., ABRAMSON,~V~.B., SCHULZ, K. F., A.ND KERKER, M., J. Phys. Chem. 64, 1157 (1960). 39. MAT~JEVIS,E., MXTHAI, K. G., AND KERKER, IV[., J. Phys. Chem. 66, 1799 (1962). 40. MATLIEVIC, E., AND KERKER, ~!/~.,J. Phys. Chem. 62, 1271 (1958). 41. TE~AK, B., MATIJ~.VId, E., AND SCHULZ,K., J. Phys. Chem. 55, 1557 (1951). 42. PaYTZ, M., Z. Anorg. Allgem. Chem. 231, 238 (1937). 43. TE~AK, B., MATIJEVI~, E., AND SCItULZ, K. F., J. Phys. Chem. 59, 769 (1955). 44. MATIJEVI~, E., SCHULZ, K. F., AND TE~AK, B., Croat. Chem. Acta 28, 81 (1956). 45. MATIJ~.VIS,E., BaOADI~UnS~,D., AND KERKER, ~[., J. Phys. Chem. 63, 1552 (1959). 46. SCnULZ, K. F., AND MATIJEVIC, E., Kolloid-Z. 168, 143 (1960). 47. PERRIN, D. D., J. Chem. Soc. 1962, 4500. 48. BIEDERMANN, G., Proc. 7th Intern. Conf. Coord. Chem., Stockholm, June 1962, p. 159. 49. HIETANEN, S., AND SILL]~N, L. G., Acta Chem. Seand. 13, 533 (1959). 50. MATIJEVI~, E., JANAUER, G. E., AND KERKER, M., J. Colloid Sei. 19, 333 (1964).