The solution of Equation 43 is easily shown to be [c(t
7O)l
1
#+%
-
=
[22/tiro]/i$%?T2/7R/Tg
(44)
This last case corresponds to the coulostatic method when the potential is a t a value iii the limiting current region foi the electrode piocess This case and it-. anal> tical applications have been di.cu\+ed b? Delahay (6) Vaiiation- of $ betneen these t n o limitiiig ca.ei can only be obtained by direct solution of Equation 40 We have u.ed the yam? numerical procedure d e w ibed for the t n o previous cases. .\I1 calculations n ere performed T$ ith 6 = 0 01 and iatios ’$z are estimated acruiate to + 0 001 Results of Calculations. From Equation 40 value. of as a function of t 7” can be calculated for any given and r n T O It nil1 be recalled that 4% is the maximum value can attain, and that fiom the definitions of T~ and T R :
+
+
+%
+
Here E , i b the initial equilibrium potential and El is the polarographic half-nave potential Thus, the parameter ( T R T ~ )in ~ Equation 40 simply defines the initial potential for the experiment in terms of the conventional polarographic M ave Typical relaxation curves calculated from Equation 40 are shonn in Figure 4 for several values of 4%and E , = El 2 . For convenience these data were nor-
malized by plotting $,‘+%. As expected, the results are for small values of identical to those calculated with Equation 42. Thus, for $%= 1 0 . 0 4 ( A E = +ll.O, n mv. at 25” C.) the values of $/ICt agree with Equation 42 to 10.001. However, even for values of $, as large as 1.0 ( A E = i 2 6 / n mv. at 25” C.), the values of $/$&in Equation 40 are only about 0.01 larger than those from Equation 42 at a given TO. Thus, in this case half-relaxation times calculated from Equation 42 are valid a t the 2% error level for potential excursions as large as 25’n mv. Actually this apparent good agreement is partly fortuitous, since the nonlinearized results (say for AE = +25 mv.) depend on both the sign of A E and on initial potential-Le., on the ratio of 7 0 and sR-whereas the linearized results depend only on the sum of relaxation constants ( T = ro TR). For example, for initial potentials anodic of El,z, the potential for large anodic excursions decays slower than Equation 41 predicts, whereas for cathodic excursions the decay is faster. The difference amounts to 15 to 2097, under some conditions. Kevertheless for many applications (especially oscillographic measurements) Equation 41 should be useful for somewhat larger potential excursions than assumed previously (9). For values of $L greater than about 5.0 the system is driven well into the limiting current region, and then the potential decay is linear with PI2, in good agreement with results predicted by Equation 44. The result that Equation 41 is useful for potential excursions greater than 2 to $$
+
3 mv. is especially important if the same situation is found for the case involving simultaneous diffusional and charge transfer relaxation. From an experimental point of view the use of potential variations larger than 1 or 2 mv. often is desirable. Thus, the possibility of using fairly large potential variations still described adequately by the relatively simple closed-form solutions, would make the coulostatic method even more attractive. LITERATURE CITED
(1) Berzins, T., Delahay, P., J . A m . Chem. Soc. 77, 6448 (1955). (2) Delahay, P., ANAL.CHEU.34, 1267
(1962). (3) Delahav. P.. J . Phvs. Chem. 66. 2204 (1962). (4) Delahay, P., “New Instrumental Methods in Electrochemistry,” p. 132, Interscience, Sew York, 1954. ( 5 ) Matsitda, H., Ayabe, Y., 2. Elektrochem. 59, 494 (1955). (6) Sicholson, R. S., unpublished work. (7) Nicholson, R. S., Shain, I., ASAL. CHEM.36. 706 11964). (8) Oldham; K. B., J : Electrochem. Soc. 107, 766 (1960). (9) Reinmuth, W. H., ANAL.CHEM.34, 1272 (1962). (10) Zbid., p. 1446. 111) Saveant. J. M.. Yianello. E.. C o m ~ t . Rend. 256,’2597 (1963). ‘ (12) Scarborough, J. B., “Numerical Mathematical Analysis,” p. 192, Johns Hopkins Press, Baltimore, 1950. (13) Smith, D. E., ANAL. CHEM. 35, 602 (1963). (14) Vielstich, W.$Delahay, P., J . A m . Chem. Soc. 79, 1874 (1957). ~
RECEIVED for review December 24, 1964. Accepted February 25, 1965. Presented in part at the Division of Analytical Chemistry, 149th Sleeting ACS, Detroit, hlich., April 1965.
Direct Potentiometric Titration of Polyethylene Glycols and Their Derivatives with Sodium Tetraphenylboron ROBERT J. LEVINS and ROBERT
M. IKEDA
Philip Morris Research Center, Richmond, Va.
b A simple, direct potentiometric titration has been developed for polyethylene glycols (PEG’s) and their derivatives, using sodium tetraphenylboron (NaTPB) as titrant in the presence of barium ions. A combined titrimetric and gravimetric procedure demonstrates that PEG’s 600 to 4000 react stoichiometrically to form complex precipitates containing 2 moles of TPB and 10.4 f 0.2 moles of ethylene oxide for each mole of barium. The approximate molecular weight of an unknown PEG may b e obtained from the infrared spectrum of its precipitate. The titration is applicable to other polyethylene oxide derivatives.
P
(PEG’s) and their derivatives have been among the more difficult classes of organic compounds to analyze. Gravimetric procedures for their determination are based on precipitation with reagents containing large anions, such as potassium ferrocyanide ( 9 ) , potassium bismuth iodide (171, and heteropoly acids (12) or sodium tetraphenylboron (KaT P E ) in the presence of barium ions ( 7 , 8, 1 1 ) . Several colorimetric methods have been investigated, most of which involve precipitation of the P E G as a colored complex which is subsequently extracted into a n organic solvent for photometric readout ( 1 , 3, I S ) . Alternatively, the precipitate is decomposed OLYETHYLEXE GLYCOLS
and one of its constituents is colorimetrically determined (4, 6). A gas chromatographic method has been reported in which the lower molecular weight PEG’S (to mol. a t . 400) have been run as their methyl ethers (2). Most of these methods require considerable manipulation and are time consuming. We desired a simple, rapid method suitable for routine use. The gravimetric procedures of Neu ( 7 ) and Seher (IO)suggested the possibility of a rapid titrimetric precipitation method for PEG’S and their derivatives. PEG’S form oxonium ions in the presence of barium, which are instantaneously precipitated by NaTPB. The complex precipitates are extremely insoluble in VOL. 37, NO. 6, M A Y 1965
e
671
aqueous or weak acetic acid medium. Therefore, during a precipitation titration the concentration of T P B ion increases rapidly in the vicinity of the equivalence point. Kirsten, Berggren, and Nilsson ( 5 ) investigated the potentiometric titration of organic bases with NaTP13. They correctly reasoned that a silver electrode would respond to changes in T P B concentration because of the low solubility of AyTPB. We have developed a simple, direct potentiometric titration of PEG’S in barium acetate buffer using NaTPB titrant. .i silver indicator electrode is used against a double junction (internal NaxO3 bridge) hg/AgCl reference electrode. Various other P E G derivatives, such as the Tweens and nonyluhenol adducts,, may” also be titrated, as may the polypropylene glycols (PPG’s). EXPERIMENTAL
Apparatus. Beckman Zeromatic p H met’er, hlodel 96, or Fisher Tit,rimeter; silver electrode, Beckman N o . 39271; double junction (fiber) si1ver:silver chloride or calomel reference electrode (add 1.0 molar sodium nitrate t,o the lower portion), Beckman No. 40452; 10-ml. buret. Reagents. Sodium tetraphenylboron ( N a T P B ) , Fisher No. S-652. An approximately 0.1 .If aqueous solution is filtered through Whatman KO. 42 filter paper using suction and a Buchner funnel. The slightly opalescent solution is stored in a clean polyethylene bottle and is stable for a t least two weeks. Barium acetate buffer, p H about 4.6, is prepared by diluting 12.0 grams of barium acetate and 6.0 ml. of glacial acetic acid to 100 ml. Standardization. For routine use the 0 . l X S a T P B titrant is directly standardized against a known weight of a n y P E G of molecular weight 600 to 4000. The reason for this range will be discussed presently. T h e NaTPB may also be titrimetrically standardized with potassium nitrate (KS03) or “Oxaditon” ( 5 ) . Figure 1 shows some typical titrat’ion curves. Procedure. Dilute 0.1 to 0.15 gram (exactly weighed) of a given P E G to about 20 ml. with distilled water (aliquots of a standard P E G solution may be used). .idd 5.0 rnl. of t’he pH 4.6 barium acetate buffer and titrate with the 0.1JI N a T P B solution using a silver indicator electrode and the double junction reference electrode. The sample solution should be stirred vigorously. 13uret and potential (millivolt) readings are taken every 0.05 ml. in the vicinity of the end point. The titration curve is plotted and the potential a t the inflection point is taken as the end point of the titration. The titration factor is calculated as gram P E G ‘ml. S a T P I I . Samples and any additional standards are then directly titrated to the established end-point potentia,l. 672
ANALYTICAL CHEMISTRY
The precautions to be observed are simple: the 1.0M N a N 0 3 bridge solution in the reference electrode should be replaced daily; the electrodes should be conditioned prior to use by making a rough titration of a P E G standard. and the electrodes should be rinsed and wiped after each titration. &Iost PEG’S contain about 1% H20; therefore, moisture determinations should be made on any PEG’s selected as
Our work with the direct potentiometric titration of PEG’S demonstrates p 1 barium 3 ions do that ~ ~ ~ and bine stoichiometrically FTith PEG'^ in the molecular weight range of 600 to 4000. w e have concluded that every 5.22 ethylene oxide units (EOc) Cornbine with 1 mole of S a T P B and 0.5 mole of barium; or 2TPB.Ba.10.4 i 0.2 EOU:
H- (OCHSCH2
[
nAa
standards. Possible interferences are organic bases, polyvinylpyrrolidone, ethylene oxide adducts, and lithjum, potassium, and ammonium ions. Anions which form insoluble silver salts will coat the silver electrode, thus poisoning it. Kirsten et al. ( 5 ) eliminated the interference of chloride and bromide ions by applying a negative voltage (-500 mv.) to the silver electrode during the titration, which repelled the halide ions, but did not interfere with the change of potential caused by the tetraphenyl borate. Since our samples were nonionic, we did not apply a negative potential to the indicator electrode. I t is also possible to remove ionic interferences, both organic and inorganic, by prior treatment of the sample with a mixed bed ion exchange resin, such as Amberlite MB-1. DISCUSSION
Seher (10) reported that the PEG’Sdo not combine stoichiometrically with KaTPB and barium, but established that the mole ratio of T P B : B a in the complex precipitates is always 2 : 1. Schonfeldt (9), however, reported that 6 units of ethylene oxide will combine with one mole of potassium ferrocyanide. Uno and Miyajima (15, 16) recently described an ingenious titrimetric procedure for nonionic ethylene oxide adducts using K a T P B as the titrant in the presence of barium ions. X pH sensitive indicator (Congo Red) is added to the sample solution, which has been adjusted to p H 3. Inclusion of the indicator in the nonionic micelle protects it from attack by hydrogen ions until precipitation of the nonionic adduct by the N a T P B destroys the micelle. The acidic color of the indicator then appears. PEG’S however, could not be directly titrated by this procedure because of the appearance of the acidic color of Congo Red when the sample solutions were adjusted to p H 3. They further concluded that 6 units of ethylene oxide combined stoichiometrically with 1 mole of N a T P B and 0.5 mole of barium for nonionics with 10 to 40 units of ethylene oxide, except for PEG’s, for which the ratio was slightly less. We have calculated from their data (by an indirect titration procedure) ratios of 6.1, 5.5, and 5.6 for PEG’S 300, 600, and 1540, respectively.
The insolubility of the TPB.Ba.PEG precipitates permits their recovery after titration. Duplicate samples of all of the readily obtainable PEG’s were titrated by the described procedure. Allowance was made for the moisture content of the PEG’S. The 0 . 1 X NaT P R was titrimetrically standardized against KNO, (Figure 1). The complex precipitates were aged overnight and filtered through tared sintered glass crucibles (fine frit), washed with 100 inl. of distilled water and then dried to constant weight in a vacuum desiccator over P205,as prescribed by Seher (10). The millimoles of T P B present in each precipitate were calculated from the titration data. The millimoles of Ba were taken to be exactly one half the millimoles of T P B ( 7 , 10, 15, 16). The respective weights of T P B and Ba in each precipitate were next calculated. The weight of PEG in the precipitate was then obtained by difference. Table I lists the pertinent data and compares the weights of PEG’s taken us. the weights of PEG’S recovered. Table I1 shows the volume of 0.099051 S a T P B consumed per gram of PEG. The end point was taken as +25 mv. for all P E G titrations. Six 0.1420gram samples of P E G 1000 gave an average titration of 6.11 + 0.01 ml. with a maximum single deviation of 0.02 ml. The combined titrimetric-gravimetric method indicated excellent agreement between the knom n and calculated weights of the various PEG’s, with the exception of PEG 1500. Table I1 shows that PEG’S 600 to 4000 gave an average ratio of milliliters of titrant to ueight of sample of 43.06 f 0.16. The ratio of PEG’S 400, 1500, and 6000, however, deviated from this value. Table I11 shows the derivation of the mole ratio of ethylene oxide to tetraphenylboron (EO/TPB) in the precipitates for the various PEG’s. I t should be noted that the average molecular weights of the PEG’S above PEG 1000 differ from their numerical designation (14). PEG 1500 is actually a 1 : l mixture of PEG’S 300 and 1540 and has an average molecular weight of 550, hllowance was made for the end groups in calculating the ethylene oxide content of the PEG’S. The EO/TPB mole
Table
I.
Titrimetric and Gravimetric Results
(0.0990M NaTPB titrant) Weight, PEG 400
NaTPB, ml TPB, mmole Ba, mmole 6 83 0 6762 0 3381 6 87 0 6801 0 3401 6 32 0 6257 0 3129 6 31 0 6247 0 3124 6 07 0 6009 0 3005 6 07 0 6009 0 3005 6 11 0 6049 0 3025 6 11 0 6049 0 3025 6.62 0.6554 0 3277 6.72 0.6653 0,3327 6.16 0.6098 0.3049 6.15 0.6089 0.3045 6.05 0,5990 0.2995 6.03 0.5970 0.2985 5,71 0,5653 0,2827 5.71 0.5653 0.2827
gram 0 1352
0 13.52 0 1472 0 1472 0 1410 0 1410 0 1420 0 1420 0,1414 0.1414 0.1417 0.1417 0.1410 0.1410 0,1400 0,1400
600 750 1000 1500 1540 4000 6000
Table II. Volume of 0 . 0 9 9 0 M NaTPB Consumed per Gram of PEG
PEG 600 750 1000 1540 4nno 400 1500 6000
RIl. TPB/gram PEG 42 93 43 05 43 03 43 47 42 84 50 67 46 82 40 79
Ml./wt.
6 6 6
6 6
6 6 5
sample 32/0 1472 07/0 1410 11/0 1420 16/0 1417 n u 0 1410 85/0 1352 62/0 1414 71/0 1400
ratio for PEG’S 600 to 4000 is 5.22 f 0.05, Again, results for PEG’s 400, 1500, and 6000 show considerable deviation. The low mole ratio of EO/TPB obtained for the P E G 400 and P E G 1500
TPB, gram 0 2158 0 2171 0 1997 0 1994 0 1918 0 1918 0 1931 0 1931 0.2092 0.2124 0.1946 0.1944 0.1912 0.1906 0.1804 0.1804
Table 111.
4v.
PEG 600 750 1000 1540 4000 400 1500 6000
mol. wt. 600 750 1000 1450 3350 400 550
6750
Ba, gram 0 0465 0 0467 0 0430 0 0429 0 0413 0 0413 0 0416 0 0416 0.0450 0.0457 0.0419 0.0418 0.0412 0,0410 0,0388 0,0388
Sample PEG, gram 0 1472 0 1410 0 1420 0 1417 0 1410 0 1352 0 1414 0 1400
EO, gram 0 1428 0 1350 0 1394 0 1399 0 1402 0 1291 0 1368 0 1400
complexes indicates that their composition differs from that of the typical complex. P E G 400 contains only 8.7 ethylene oxide units per mole, and P E G 300 (a major constituent of P E G 1500) only 6.4 ethylene oxide units per mole.
0
gram
0 1325
0 1314 0 1467 0 1462 0 1416 0 1416 0 1428 0 1420 0 1346 0.1296 0.1422 0 1423 0.1396 0.1431 0.1417 0.1396
EO, mmole 3 25 3 07 3 17 3 18 3 19 2 93 3 11 3 18
TPB, mmole (from
titration) 0 6252 0 6009 0 6049 0 6094 0 5980 0 6782 0 6604 0 5653
EO/TPB 5 20 5 11
5 5 5 4 4
24 22 33 32 71 5 63
This suggests that the spatial arrangement of the atoms within the ciyital lattice of the complex demand> a minimum chain length of ethylene oxide units (approximately 10.4) to i a t i d y a preferred structure. Although each of the complex precipitates obtained from PEG’s 600 to 4000 contains the same weight per cent of P E G , each contains a different mole ratio of PEG. - h o t h e r way of simply
n
40! 30
1.0
-
0.8
-
0.6-
IO00 -50-
PEG by diff.,
Derivation of Mole Ratio of EO/TPB
0
40-
+
TPB Ba, gram Ppt., gram 0 2623 0 3948 0 2638 0 3952 0 2427 0 3894 0 2423 0 3885 0 2331 0 3747 0 2331 0 3747 0 2347 0 3775 0 2347 0 3767 0,2542 0 3888 0.2581 0.3877 0.2365 0.3787 0.2362 0,3785 0.2324 0.3720 0.2316 0.3747 0.2192 0.3609 0.2192 0.3588
0 0
0
e
8,
6
e
0.4-
a
\
2 TPB.Ba. nPEG
0
0.2-
1,000
2,000 3,000 4,000 5.000
6,000 7.000
MOLECULAR WEIGHT OF PARENT PEG Figure 2. Plot of n vs. molecular weight of parent PEG for 2TPB.Ba.nPEG complex precipitates VOL. 37, NO. 6, M A Y 1965
673
10,000-
0 W
0
E : i3
ABSORBANCE
I ,1
,
,
,
,
3
4
5
6
1
I
,
.
.
.
.
"
I
. .
1 0 11 1 2 1 3 14 1 5 MICRONS 8
9
Figure 3. Infrared spectra of NaTPB (Nujol), PEG 1000, and 2TPB.Ba.nPEG 1000 precipitate (Nujol)
presenting the structure of these precipitates is ~ T P B . U ~ T L P E where G, n is a function of molecular weight or n =
moles P E G moles Ba
~-
Figure 2 shows a plot of the mole ratio of P E G in the complex precipitates us. the molecular weight of the parent
PEG. This variance in the mole ratios of the PEG'S in the precipitates gave rise to speculation that the differences might be reflected in their infrared spectra. Figure 3 shows the infrared spectra of PEG 1000, NaTPB, and the 2TPB.Ba. aPEG complex. Similar spectra were obtained for the other precipitates. The bands a t 9.2 and 14.1 microns were assigned to the PEG ether stretching frequency and to the substituted phenyl ringb of TPI3, respectively. The ratio of the absorbance of these two bands nab plotted against the log of the molecular w i g h t of the parent PEG, as shown in Figure 4. This curve dem-
Table IV. Molar Solubilities and Melting (Decomposition) Poihts of 2TPB.Ba.nPEG Precipitates Melting Jlolar point, PEG solitbility OC. 400 600 750 1000 1500 1540
4000
6000
674
1 6 X 10-6 1 0 X 10-6 1 6 X 10-5
2 4 X 10-6 2 7 X 10-6
219-224 230-233 244-247 239-242 239-244 237-240 240-243 243-246
ANALYTICAL CHEMISTRY
RATIO
Figure 4. Plot of absorbance ratio 9.2 microns/ 14.1 microns vs. log molecular weight of parent PEG for various 2TPBsBaqnPEG precipitates
onstrates that the approximate molecular weight of a P E G can be obtained from the infrared spectrum of its complex precipitate. The molecular weight cannot be obtained from Figure 2, since the curve shown there demands prior knowledge of the number of moles of PEG, and 'hence of its molecular weight. The molar solubilities of the complex precipitates having the empirical formula 2TPB.IIa.10.4 + 0.2 EOU (PEG'S 600 to 4000) were determined from the absorbance of their saturated aqueous solutions a t 265 mk. Each mole of the dissolved complex yields two moles of TPB. Standard solutions of N a T P B were used to make a calibration curve of absorbance us. concentration. The solubilities are shown in Table IV. Sinsheimer and Smith (11) have shown that the melting points of the T P B complexes of amines can be useful in their identification. Therefore, melting points of the P E G complexes were obtained on a Fisher-John's apparatus. We had expected that the precipitates would melt or decompose a t quite low temperatures, since Seher (10) states that the complex salts must be dried a t room temperature because of their loa melting points. This is the reason the precipitates were dried in a vacuum desiccator over P,Os. S e u ( 7 ) also states that the 2TPB.IIa~zTween precipitates should be dried below 50' C. in vacuo. However, we found that all of the PEG complex precipitates melted (1% ith decomposition) between 219' and 246' C., with no sign of any physical change below 200" C. The melting points of the precipitates are shown in Table IV. A DTA curve was run on the P E G 1000 complex precipitate (under N, atmosphere). I t melted sharply a t 252' C. with no sign of prior decomposition.
CONCLUSIONS
The potent'iometric titration procedure for determining PEG'S has several advantages over other methods. It is simple, fast, and direct. It is carried out in aqueous solution, and PEG'S can be eluted from most base materials with hot water. Other glycols and their esters-e.g., glycerin, propylene glycol, t'riethylene glycol, and triacetin-do not interfere. In routine use, no isolation of a precipitate or colored complex is required. Well defined titration curves may be obtained with 0.025&VKaTPB (about 2 mv. per 0.01 ml. in the vicinity of the inflection point). Reductions in the volumes of solutions would permit semimicro determinations, as demonstrated by Kirsten, et al. ( 5 ) . Infrared spectra of the isolated precipitates can give useful information on the parent materials. The Tween precipitates, for example, give typical spectra with a small but distinct carbonyl band. The precipitation step may also be regarded as a separation and isolation technique. Finally, the composition ,of the complex P E G precipit'at'es has been established to be 2TPB.Ba.10.4 + 0.2 EOU. The precise composition of other ethylene oxide complex precipitates was not' investigated. LITERATURE CITED
( 1 ) Brown, E. G., Hayes, T. J., Analyst 80, 755 (1955). ( 2 ) Celades, R., Paquot, C., Rec. Franc. Corps Gras 9 , 145 (1962). ( 3 ) Coppini, D., Cameroni, R., Boll. Chzm. Farm. 92, 363 (1953). ( 4 ) Ken-ichi Hattori, et al., J . Chem. SOC. Japan, Ind. Chem. Sect. 64, 1195 (1961). ( 5 ) Kirsten, Vi'. J., Berggren, A., iiilsson, K., ANAL.CHEM.30, 237 (1958). ( 6 ) Morgan, D. J., Analyst 87, 233 (1962).
(7) Neu, R., Artnezmzttel Forsch. 9, 585 (1959). (8) Seii, R., Fette, Sezfen, Anstrzchmztted 61, 980 (1959). ( 9 ) Schonfeldt, N,,. J . A m . Oil Chemists’ Soc. 32, 77 (1955). (10) Seher, 4., Fette, Seifen, Anstrichmittel 63, 617 (1961).
(11) Sinsheimer, J. E., Smith, E., J . Pharm. Sci. 52, 1080 (1963). (12) Smith, W. B., ilnalyst 84, 77 (1959).
(13) Stevenson, D. G., Ibid., 79, 504 (1954). (14) Cnion Carbide Chemicals Co., 270 Park Avenue, Kew York 17, N. Y., booklet “Carbowax Polyethylene Glycols” (1960).
RECEIVED for review January 21, 1965. Accepted March 17, 1965.
Chronopotentiometric Evidence for the Formation of Europium(ll1) Acetate Complexes DANIEL J. MACERO, HARVEY B. HERMAN,’ and ALEXANDER J. DUKAT Department of Chemistry, Syracuse University, Syracuse, N. Y. The reduction of Eu(lll) to Eu(ll) was studied by voltammetry at constant current in sodium acetate and acetic acid solutions at constant ionic strength. The pH was kept constant at 4.5. Concentration vs. the square root of transition time plots a t a current density of 0 . 0 2 8 2 ma. per sq. cm. were linear in accordance with the Sand equation. A diffusion coefficient of 0 . 4 0 X 10-5 sq. cm. per second was evaluated for Eu(lll) in 0.1M acetate solution at 2 5 ” C. The product, i ~ ~ j ~ was / C , constant over a range of current densities and europium(lll) concentrations. The electrode reaction for the reduction of Eu(ll1) to Eu(ll) in the acetate system was found to be reversible from current reversal studies and potential vs. time plots. Evidence is given for the presence of the acetate complexes, EU(OAC)~+ and Eu(OAc)z+. The values of the corresponding stability constants were calculated to b e 3 2 4 and 20.4.
S
of the voltammetric behavior of europium(II1) in aqueous solution have been restricted t o polarographic studies with the dropping mercury electrode. In all of these the work was done wit,h unbuffered solutions of either neutral or acidic pH. Only a single wave corresponding to the reduction of Eu(II1) to Eu(I1) was obtained by Soddack and I3ruckl (11) and by Laitinen and Taebel (8). Both groups of investigators concluded that the electrode reaction was reversible or nearly reversible. Two waves, interpreted as reduction to Eu(I1) and Eu(O), respectively, were reported by Holleck (6) and rnore recently by Misumi and Ide (IO). Onstott (12) reported the polarographic behavior of europium in EDTA solutions. Gierst and Cornelissen ( 5 ) studied the influence of various supporting electrolytes on TUIIIES
Present address, Department of Chemistry, University of Georgia, Athens, Ga.
1 32 IO
the shape of the polarographic waves for the Eu(I1)-Eu(II1) oxidation. hnderson and Nacero ( 1 ) investigated the chronopotentionietric behavior of europium(II1) in perchlorate solutions and calculated a value of -0.35 volt us. N.H.E. for the formal reduction potential at 25” C. of the Eu(II1)-Eu(I1) couple in 1 M sodium perchlorateperchloric acid (pH 2.08) solution. I n this paper the reduction of europium(II1) to europium(I1) in acetate solutions is examined by the technique of chronopotentiometry. Only the reduction step of Eu(II1) to Eu(I1) can be studied with this technique, however, since the evolution of hydrogen occurs a t a potential well in advance of that corresponding to the reduction of Eu(I1) to Eu(0). EXPERIMENTAL Apparatus. The constant current was supplied by a Heathliit Model PS-4 regulated power supply. T h e current was determined prior to each r u n by measuring the voltage drop across a General Radio Type 500-H precision resistance in series with the electrolysis cell. The electrolysis cell was adapted from a design by Delahay and Mattax ( 3 ) . The pool potential was measured against a saturated KCl, silver-silver chloride reference electrode. A platinum gauze electrode served as the anode and was separated from the cathode compartment by a medium porosity fritted glass disk. Solutions were deaerated by bubbling purified nitrogen saturated with water vapor through the solution for 5 to 10 minutes before a run and over the solution during the course of the run. The solution was allowed to come to rest for 1 to 2 minutes after outgassing and before making a run. The entire electrolysis cell was mounted on a heavy stand placed on a partially inflated inner tube (16) and carefully leveled before each determination. All measurements were made a t 25.0” i 0 . l 0 c. The output from the electrolysis cell was fed into a Beckman Zeromatic pH
meter and this in turn to a Minneapolis Honeywell Type Y153X recorder with 2.5-mv. full scale sensitivity, 0.5-second full scan and chart speeds of 8 inches per minute and 4 inches per second (with a special gear train). The p H meter recorder output was shunted with a variable resistance to expand the voltage axis of the Honeywell recorder. Potential measurements were reproducible to 1 2 mv. A Beckman Model G pH meter [Vas used to determine p H values. Reagents. All chemicals with the exception of the europium were analytical reagent grade. Stock europium solutions were prepared by dissolving 99.8YG pure europium oxide, Lindsay Chemical Division (American Potash and Chemical Corp.) in a small quantity of nitric acid and diluting to volume. This solution was analyzed volumetrically by an EDTA titration (17). For each chronopotentiometric determination an aliquot of stock europium nitrate solution was mixed with sufficient sodium acetate solution, acetic acid, and sodium perchlorate solution to provide a buffer system approximately 1 J Z in ionic strength. RESULTS AND DISCUSSION Reversibility of Electrode Reaction. A typical chronopotentiogram with
reversal of current is shown in Figure 1. This was taken with a 1.18 x lO-3M E U ( S O ~ )solution ~ containing 0.1M sodium acetate and 0.L21 acetic acid with sufficient sodium perchlorate added to give an ionic strength of 1M. The p H of the solution was 4.5 and the current density was 0.0282 ma. per sq. em. The quarter-wave potential of the cathodic portion occurs a t -0.705 volt us. S.C.E. and the “quarter-wave” [actually the 0.215 wave ( I S ) ] potential of the anodic curve a t -0.707 volt us. S.C.E. This excellent agreement between the forward and reverse quarterwave potentials is strong evidence for the conclusion that the electrode reaction is reversible. Similar agreement between the cathodic and anodic quarVOL. 37, NO. 6, M A Y 1965
675
