J O U R N A L OF T H E AMERICAN CHEMICAL SOCIETY (Registered in U. S. Patent Office)
(Copyright, 1965, by t h e American Chemical Society)
OCTOBER 27, 1955
VOLUME77
N U M B E 20 R
PHYSICAL A N D INORGANIC CHEMISTRY [CONTRIBUTIONFROM
THE
NOYES CHEMICAL LABORATORY, UNIVERSITY
O F ILLINOIS]
Polarographic Study of Tris-(ethylenediamine)-cobalt(111) Ion BY H. A. LAITINEN AND M. W. G R I E B ~ RECEIVED APRIL20. 1955 The polarographic reduction of tris-(ethylenediamine)-cobalt(111) ion is shown to be reversible in the presence of excess ethylenediamine. The anodic half-wave potential for the corresponding cobalt( 11) complex is identical with that of the cathodic process. A study of the formation of ion pairs by tris-(ethylenediamine)-cobalt(111) ion with certain anions was made, and the magnitudes of the equilibrium constants for the dissociation reactions were estimated.
Polarographic reduction curves of tris-(ethylene- characteristic of the reversible potential of the trisdiamine) -cobalt(III) ion were observed by Herda2 (ethylenediamine)-cobalt(II1)-(11) system.' I n the present investigation, the conditions for in potassium nitrate solutions in the presence of gelatin and octyl alcohol, and by Willis, Friend and reversible reduction were studied, and ion-pair Melior3 in potassium sulfate and sodium acetate formation reactions were quantitatively evaluated. solutions. I n neither case was excess ethylenediExperimental amine present and, as will be seen below, the conApparatus and Procedure.-A Sargent Model XXI Poditions for a reversible reduction were therefore not ldrograph, a Leeds and Northrup Model E Electrochemofulfilled. Bjerrum4 evaluated the consecutive form- graph, and a Sargent Model I11 Manual Polarograph were ation constants of the complexes of cobalt(I1) ion used, together with suitable external potentiometers. For with ethylenediamine in 1 N potassium chloride, manual observations, the mean current was measured by a and the potential was determined by an exand evaluated the formal potential of the tris- galvanometer ternal potentiometer, corrected for the potential drop due (ethylenediamine) - cobalt(II1)-tris - (ethylenedia - to the average resistance of the cell. Alternatively, manual mine)-cobalt(I1) system in 1 N potassium chloride measurements were made with the recording polarographs, a t 30" a t -0.259 volt (us. the normal hydrogen taking as average currents the mean values of the small oscillations of the current recorder operated a t maximum electrode). Concurrently with and subsequent to damping. Polarograms were recorded with the Sargent the present investigation, Kivalo5 has studied the Model XXI Polarograph, using a span voltage of 0.5 v . , effect of the instability of the tris-(ethylenediam- making two potential measurements with a potentiometer, ine) -cobalt(II) ion under certain experimental con- after stopping the instrument, a t points on the wave before after the half-wave potential. Current values were ditions on the cathodic behavior of the cobalt(II1) and taken corresponding to the envelope of the tops of the fluctuacomplex. The anomalous faradaic impedance of an tions of the recorder, operated without damping. Potenelectrode caused by the adsorption of both electro- tial values were corrected for iR drop, using the mininium active species a t the equilibrium potential has re- value of cell resistance a t the moment of the fall of the drop, with a 1000 cycle a.c. conductivity bridge. cently been described.6 The reduction of hexam- measured The solution was introduced into both halves of an Hminecobalt(II1) ion in the presence of excess ethyl- type cell with a fine porosity fritted diaphragm, using no enediamine has been shown to occur a t a potential agar plug. The side arm of a saturated calomel electrode (1) Abstracted from the P h D. Thesis of M . W. Grieb, 1953. (2) M. G. Herda, X S . Thesis, University of Illinois, 1943.
(3) J. B. Willis, J. A. Friend and D . P . Mellor, THISJ O U R N A L , 67, 1680 (1945). (4) J. Bjerrum, "Metal Ammine Formation in Aqueous Solution," P . Haase and Son, Copenhagen, 1941. ( 5 ) P. Kivalo, Ph.D. Thesis, .&bo Akademi. .&bo. Finland, 1954; THIS J O U R X A L , 77, 2078 (1955). (6) H. A. Laitinen and J. ti. B. Iiaudles, Trtrns. Purodoy .Sue., 51, 54 (1055),
was introduced into one side of the cell, in which the solution was maintained a t a level below that of the solution to be electrolyzed in the other side. The latter solution was deaerated before the dropping mercury electrode was introduced. The cell and reference electrode were maintained a t 25 rt 0.02" by a water thermostat. Two different capillaries were used. The characteristics of capillary I are given in Table 11. Capillary 11, with a drop time of 3.52 5ec. and m a s rate of floir of 3.12 m g .
scc.-l, measured on open circuit i n air-saturated 0.1 h ' potassium chloride, was used in a few experiments for which only potential data are reported. Materials.8-Tris-(ethylenediamine)-cobalt(III) chloride mas recrystallized from concentrated aqueous solution b\addition of concentrated HCl and ethanol, and was dried for two hours a t 100'. Stock solutions (0.05 M ) were stored over periods up to three months with no evidence of decomposition. Tris-(ethylenediamine)-cobalt(II1) perchlorate was prcpared from the sulfate, previously recrystallized by a procedure similar to that used for the chloride, by metathesis with a slight excess of barium perchlorate solution acidified vith a drop of perchloric acid, followed by filtration of the barium sulfate. It was purified by repeated recrystaliization and washing with small portions of distilled water, until the washings showed no turbidity upon addition of sulfuric :tcid and were nearly neutral t o litmus. The salt, unlike the chloride and sulfate, is soluble in ethanol and only moderately soluble in water. A stock solution, approximately 0.5 M of cobalt(I1) rhloride was made from Maliinckrodt Analytical Reagent grade hexahydrate product. Seventy per cent. ethylenediamine-water azeotrope (Union Carbon and Carbide) was distilled after addition of :L few pellets of sodium hydroxide, a middle constant boiling fraction being collected. -4 1 M stock solution was standxrdized by potentiometric titration with standard acid. Upon prolonged standing, the stock solutions showed evidence of oxidation in a siight rise in the diffusion current a t -0.55 v. v s . S.C.E.
Results and Discussion Effect of Ethylenediamine Concentration.-The cathodic curves of millimolar tris- (ethylenediamine)-cobaltiIII) chloride in 0.1 -11 potassium chloride and various concentrations of ethylenediamine were studied using capillary I, with a merfury column height of 70.0 cm., and recording the curves with the Sargent hlodel XXI polarograph. T o determine the anodic curves of the cobalt(I1) complex, a residual current curve was first recorded. Then to 25 nil. of air-free solution was added one drop of cobalt(I1) chloride stock solution, with passage of purified nitrogen through the solution. This procedure yielded an approximately millimolar solution of tris-(ethylenediamine)-cobalt(II) ion. The trace of oxygen introduced with the drop of cobalt(I1) solution caused sufficient oxidation of the cobalt(1I) complex to be detected, but conipetisation was made in the calculation by appropriate adjustment of the residual current. The half-wave potentials and reciprocal slope values obtained from the logarithmic plots are given in Table I. The constancy of the half-wave potential a t ethylenediamine concentrations greater than 0.03 .11indicates that under these conditions both coIialt(II1'1 arid cobalt(I1) are coordinated with ail equal nuniber of ethylenediamine molecules, r~liz., three. The reciprocal slopes show satisfactory agreement with the theoretical value for a reversible reaction a t least down to 0.005 -21ethylenediamine. I n the absence of excess ethylenediamine, the behavior is not reversible.s In each case, except that of the cobalt(I1) solution in 0.004 II ethylenediamine, the plot of log [(id)c - i] 1 [i - ( i d ) a ] 21s. E gave very good straight lines. The excepted case exhibited only slight deviations a t potentials more negative than that of the half-wave. Diffusion currents in the various solu(81 'l'lir samples c v i n j i l e s ciiinpound.: iiscd in thi-. \ t i l i l y ~ r i i e i ~ i i s l~y~ ~ ~ > j >l i )l i Iw' rl~ ~ f e s w r C . 13;iil:ir, J r
M CI-C
TABLE1 \'ARIA4TIUX II.ITH C O N C E K T R A T I O S O F ETH1L E X E U I A h I I X E OF
HALF-\I.AVE P O T E N T I A L S A N D K E V E R S I R I L I T Y \'AI,UBS OF [ C o ena] I o r . ~ N D[Co en,] + + Ion- FORMED IX THE SUPPORTING ELECTROLYTE -++
Cuucn. e n .l I ii . .j
.4 7
.i
,1 1 ri7
.06 .03 , 02 .01 ,007
,005 ,004 .003 . 00 I
- -- Anodic wave---
----Cathodic wave--E ' I , , v. i'S . Reciprocal S.C.E. slope, Y
-0.459 , 461 - ,456 - ,457 - 458 - 4.57
E , ( 2 s v. us. 5.C.E.
-0.457 - ,459 - ,487 - . -157
0,0604 ,0587 ,0564 ,057i i
,457 - ,456 - ,434 - ,462 - .4*50 ..... - ,444 - ,430 -
Reciprocal
0602 . O;i81)
-
,456
--
,455
0570 058.5 ,0572 .0.581
-
,45.i
- ,4.% - ,453 -
450
- ,446 - ,434 - .429
0584 0878 .... . (1.575 ,0626
..... .....
tions were constant within the accuracy of the method. Cobalt(I1) solutions in the two lowest concentrations of ethylenediamine, for which data are not reported, showed progressively much reduced diffusion currents, indicating t h a t all available cobalt was not complexed. Diffusion Current Constants.--Extensive dat'i were obtained to verify that the limiting currents for both the cathodic and anodic processes were diffusion controlled. The average diffusion currents and drop times were measured a t constant potential on the plateau of the waves, with the mercury reservoir a t various heights. T h e mass rate of flow of mercury was determined under similar conditions. Appropriate back-pressure corrections were made.g A few typical data are presented in Table 11. The anodic diffusion current constants could not be evaluated precisely because of uncertainty in the absolute concentration of cobalt(I1) complex, but the data showed similar consistency at various values of the height of reservoir. TABLE I1 CONSTANCY OF id/hl/? A N D I (a1 1.00mill[Coenr]CIJin0.1MKC1,O.l M e n ; capillary I . (1)) 2.50 m J I [Co ena]Cla in 2.0 M KCI, 0.1 M en; cnpillary I . I, pa.
',
Sec.':2/
I!, cm.
mg. 'sec.
sec.
(a) 70.0 70.0
1.690 1.690
4.41 4 34
2.78 2.79
30 .n
1 . ion 6.21 1 440" 5.12 7u. 0 1 690 4.35 3.76 80.0 1.937" 00.0 3.35 2.183" IO0 . (1 2 -150 3 01 a \.ulues 111- ititerpcil'ttioii.
5 . 87
CjO.(l
0.4.i 0.96
(1,)
111,
m.\..I
id,
mg.2!J
pa.
7.16 7.93 8 36
1.530 1.539
0.8433 ,8430 ,8414
,8424 ,5433 %2R
1.543 1.541 1.536 1.539 1.541 1 ,539
Effect of Supporting Electrolyte.-Using a 1.OO millimolar solution of tris-(ethylenediamine)-cobalt(II1) chloride, in 0.1 31 ethylenediamine, and capillary I a t a mercury column height of 70.0 cm., the effect of various concentrations of supporting electrolytes on the half-wave potentials and diffusion current constants was determined. .I11 half-wave potentials were corrected for potential drop due to cell resistance. Anodic half-wave potentials were found to coincide with those of the cathodic waves within an experimental error of about 1 mv. and are not separately reported. a. Sodium Perchlorate.-From polarograms recorded with the Sargent Model X X I polarograph, the half-wave potentials were found to remain essentially constant a t an average value of -0.441 v. US. S.C.E. as the concentration of sodium perchlorate was varied from 0.1 to 4.0 M . However, precipitation of potassium perchlorate occurred a t the junction with the saturated KC1-agar salt bridge, causing a moderate increase in cell resistance and lowering the reproducibility of the potentials. The apparent constancy of half-wave potential with sodium perchlorate concentration has no particular theoretical significance, having arisen presumably through cancellation of relatively small junction potentials and activity coefficient effects. The experiment was repeated with capillary I1 and a solution of tris-(ethylenediamine)-cobalt(II1) perchlorate, estimated by comparison of diffusion currents to be 0.5 millimolar, in 0.03 sodium perchlorate. The agar plug was protected by saturated potassium chloride in the bridge side of the Hcell. No precipitation of potassium perchlorate was observed, and the cell resistance was normal. The half-wave potential thus obtained was - 0.4414 v. vs. S.C.E., and was used to calculate the “standused in subsequent ard” half-wave potential (El,20)s calculations. b. Potassium Chloride.-The half-wave potential data in Table I11 were obtained with the Sargent Model I11 manual polarograph.
diamine was ignored. To estimate values of the junction potentials, the formula of MacInnesll was used, with the e.m.f. data of Shedlovsky and MacInnes.I2 When plotted against logarithm of innlar concentration, the junction potentials gave a smooth, nearly linear plot. T o evaluate the formation constants of ion-pair or higher ion-cluster complexes, the method of DeFord and HumeI3 was used, assuming that the divalent complex ion does not form ion-pairs. The functions Fo(X),Fl(X) and F2(X) are defined as by DeFord and Hume, but regarding the tris-(ethylenediamine) -cobalt(III) ion as the central ion. ‘To calculate these functions, i t is necessary to adopt a value of the half-wave potential of the system in the absence of complexing anions. The best reference value seems to be t h a t measured in 0.03 M sodium perchlorate, corrected for liquid junction potential and for activity coefficient ratio of the trivalent and divalent cobalt complex ions. The junction potential is estimated to be 2 . 5 mv., due principally to diffusion of the saturated potassium chloride, giving -0.4839 v. vs. S.C.E. for the corrected potential. Estimating the activity coefficients by a formula used by Kielland,I4 taking an ionic radius of 3.68 A. for each ion, we estimate the “standard” half-wave potential a t unit activity ratio to be -0.420 v. vs. S.C.E. The figure 3.68 8. was obtained from limiting equivalent ionic conductance of the tris-(ethylenediamine)-cobalt(II1) ion as determined by Jenkins and Monk,lBand is consistent with crystallographic interatomic distances in both complexes.16--19 The values of Fo(X) and F,(X) are plotted in Fig. 1. Extrapolation of the Fl(X) to zero concentration of supporting electrolyte gives the value 28 for the formation constant of the [Co ens]Cl++ ion pair. The dissociation constant is the reciprocal, 0.036. The slight upward trend of Fl(X) a t high concentration indicates the beginning of association of two chloride ions with each cation. The marked upward trend of Fl(X) values a t the lowest concenTABLE I11 trations may be due to ignoring the activity coefANALYSISOF DEPENDENCE OF E l / , ON POTASSIUM CHLORIDE ficient ratio of the ion-pair and the divalent cobalt CONCENTRATIONS IN SOLUTIONS CONTAINING1.OO m-lf complex in the calculation (it was assumed that the [Co en31 Cla A N D 0.1 M en. ion pair behaves more nearly as a divalent ion than KCl, El/,, E ~ , cor., AEil2, as a trivalent one, so t h a t this ratio could be negllrI uKC1 v. v s . S.C.E. v. v. us. S.C.E. v. lected). Over-correction of activity coefficients in 0 0 -0.420 estimating the “standard” half-wave potential would 0.01 0.009 -0.436 0.0031 - ,433 0.013 also cause such an effect. For example, if the value ,025 ,021 - ,442 ,0027 - ,439 ,019 -0.430 v. vs. S.C.E. is adopted for (El,,:)s, the .05 ,041 - .448 ,0024 - ,446 ,026 Fl(X) values a t low concentrations are brought .10 ,077 - ,456 ,0020 - ,454 ,034 .20 ,144 - ,4652 ,0017 - .4635 ,0435 much closer to the extrapolated portion of the curve, .40 .269 - ,4765 ,0013 - ,4752 ,0552 the formation constant is estimated to be 18, and .70 ,447 - ,4880 ,0010 - ,4870 ,0670 the dissociation constant to be 0.056. 1.0 2.0 3.0 4.0
,623 1.225 1.899 2.686
-
,4956 ,5125 ,5266 ,5390
.0008
.0004 ,0002 ,0000
- ,4948 - ,5121 - .5264 - ,5390
,0748
,0921 ,1064 ,1190
The activities of potassium chloride in the various solutions were calculated from an equation expressing the mean activity coefficients as reported by Shedlovsky.lo The effect of the 0.1 Methylene(101 T . Shedlovsky. T H I SJ O U R N A L . 72, 3080 11950).
D. A. MacInnes, ibid., 37, 2301 (1915). T. Shedlovsky and D. A. MacInnes, ibid , 59, 503 (1937) D. D. DeFord and D. N. Hume. ibid., 73, 5331 (1951). J. Kielland, ibid., 69, 1675 (1937). I. L. Jenkins and C. B . Monk, J . Chcm. Soc., 68 (1951). M. Atoji and T. Watanabe, J . Chem. P h y s . , 20, 104.5 (1952). D. P. Mellor and R. J. Goldacre, J . Proc. R o y . Soc. S.S. Wales, 78, 233 (1940). (18) L. C. Pauling, “ T h e N a t u r e of t h e Chemical Bond,” 2nd. ed.. Cornell University Press, I t h a c a , S . Y . , 1915, pp. 179, 182. (19) T. Watanabe and J. a t o j i , Scieptcc ( J a p a n ) , 20, 422 (1950): C. A , , 46, 997Zb (1951). (11) (12) 113) (14) (15) (16) (17)
I
8OC
TABLEI \ ' EFFECTOF SasSOa COYCENTRATIOS ON HALF-KAVE POTENTIAL
1.00 m X [ C o ens] cla and 0.1 k f en
.i
P
Molar Concn. NaqSO4 Sod-
0,Ol ,023 0.;
1 ,7>
0.009.5 0230 ,0442 0844 ,183
. I )
1.o 2 ,5 ,
~
y\c( A C T I V I T Y . Fig 1.-Variation
of ,F(x) with activity of KCl.
These values are in reasonably good agreement with the value of 0.019 for the dissociation constant, calculated from conductance data by Jenkins and who regard it as approximate to an order of magnitude. From the present data, i t appears fairly certain that 0.036 is a minimum value for this constant. Sodium Sulfate.-The half-wave potentials in Table IV were obtained with the Sargent Xodel XXI recording polarograph operated by the manual method previously described. Although it is seen that sulfate ion causes a very large shift of half-wave potential, it has not been possible to attempt an accurate calculation of its formation constant. No suitable method is known to correct for the junction potential between saturated potassium chloride and concentrated sodium sulfate solutions. The effect on potential of the activity coefficients of the ion pair and the reduced ion cannot be expected to cancel, because they differ in charge. The activity of the sulfate ion in concentrated solutions remains in doubt.
El,%, v. L'S.
S.C.E.
-0.420 - ,477 - ,493 - ,505 - .SlS
- ,537 ,557 - ,580 - ,630
-
AEl/,. v.
0.057 073 ,085 .nR8
,117 ,137 ,160 ,210
Fp
(X)
F i (X)
9 . 1 9 862.2 1 7 . 13 701 ..'I
3i.37 595
.i
G . 3 1 52.5.0 94.91 315.9 206.6 505.8 3540
As an approximation to the latter problem the F,(Xj values in Table IV have been calculated on the basis of free sulfate ion concentrations found in data of Righellato and Davies") for 18". These authors concluded from conductance data that sodium and sulfate ions in solution undergo ionic association. Because in the region for which the conductance data were available, the F1(X) values decrease rather than increase with increasing concentration it would not be meaningful to attempt an extrapolation. If, by analogy with the case of the chloride, the first values are disregarded, the formation constant may be estimated as about 5 2 5 . The corresponding This value is dissociation constant is 1.9 X in only fair agreement with the value 3.55 X lo-: calculated by Jenkins and Monk15 from conductance data. Sodium Nitrate.-Qualitative evidence for a weak ion-pair complex with nitrate ions was obtained from the fact t h a t the half-wave potential varied from -0.451 t o -0.52-1 v. Z'S. S.C.E. as the concentration of sodium nitrate was varied from 0.1 to 4.0 Jf in the presence of 0.1 AZI ethylenediamine. From the fact that the shift of potential is smaller than with chloride ion, we may infer that the nitrate ion forms a less stable ion pair. Acknowledgment.-This investigation was supported in part by the Office of Naval Research under Contract Xo. N6-orc-07136. L-RBASA,ILLINOIS
