SPECIESOF COBALT(II) IK ACETICACID
Kov., 196 i
1993
SPECIES OF COBALT(I1) IN ACETIC ACID. PART 11. COBALT(I1) I N THE PRESEIVCE OF LITHIUM BROMIDE, LITHIUM CHLORIDE ,4ND AMJIONIURI THIOCYANATE BY P. J. PROLL AND L. H. SUTCLIFFE Department of Inorganic and Physical Chemistry, University of Liverpool, Liverpool, England Reeemed April 86,1061
The effrrt of the addition of inorganic bromides, chlorides and thiocyanates to solutions of cobalt(I1) in acetic acid has been studied in detail, over the temperature range 25.0 to 64.0'. From ion migration and spectrophotometric experiments the ionic species present are postulated to be Co(X)+, Co(X)*, Co(Xh- and C O ( X ) ~ ~&.here -, X represents the bromide, chloride or thiocyanate, the latter two species being favored by the addition of the corresponding inorganic salts. The equilibrium constants K2 and K? for the formation of the Co(X),- and the Co(X),Z- have been evaluated for the bromide and chloride a t 25", and the product of these equilibrium constants for the thiocyanate has been found. An attempt has been made i o calrulate the oscillator strenzths of the Co(X),Z- ions; for the chloride ion the agreement with previous workers is reasonable.
Introduction In a prrvious paper from this Laboratory' it has been shown that the heating of, or the addition of inorganic acetates to, cobaltous acetate in anhydrous acetic acid converts the pink octahedral form Co(OAc)2.4HOAc into the blue tetrahedral form CO(OAC)~~-, the latter species having an absorption maximum at 56.5 mp and an extinction coefficient of about 1700 A1-l cm.-1, and the former species a maximum a t 526 mp with an extinction coefficient of 17.0 M-' cm.-l. Two equilibria were found to be awociated with this conversion, namely
+ OAC( 2 0 ( 0 4 r ) ~ -+ OAc-
K*
C0(04c)?
CO(OAC)~K3
CO(OA~)~~-
form may lead to erroneous conclusions when applied to a system in which two coordination numbers are exhibited by the central metal atom. Much of the work reported in the literature on similar systems has been done using this continuous variation method and is therefore unsatisfactory. Various alcohols have been employed as solvents but since tetrahedral complexes of cobalt(I1) are favored by a low dielectric constant, acetic acid is a better solvent for the study of these complexes. Apart from postulating the species present little attention has been given so far to the evaluation of the equilibrium constants involved in these systems. The present paper is concerned with the bromide, chloride and thiocyanate complexes of cobalt(I1) in anhydrous acetic acid and the evaluation of the equilibrium constants concerned.
Tt also was possible to estimate the values of the Experimental constants K z and K3 from spectrophotometric The necessity for obtaining completely anhydroufi reagents measurements. The importance of these equilibria has been shown to be very important in our previous paper.' in both catalytic and non-catalytic reactions was The preparation of anhydrous acetic acid and cobaltous acetate already have been described.' The precautions demonstrated and it was concluded that C O ( O A ~ ) ~ ~ is probably the most reactive species. Cobaltous taken in the preparation of the other reagents are as followP. Cobaltous Bromide.-The British Drug Houses (B.D.H.), bromide or cobaltous acetate with added lithium product was recrystallized from purified acetic acid and then bromide in anhydrous acetic acid also is used as pumped under a vacuum for about a week a t 100" until the an autoxidation ~ a t a l y s t ,but ~ , ~little attention has bright green anhydrous product was obtained. Cobaltous Chloride.-The Oakes and Eddon product Fwn paid to the role of the catalyst, hence it was decided to make a detailed spectrophotometric was recrystallized from purified acetic acid and then pumped a vacuum for about a week a t 100' until the very investigation of this system. At the same time the under pale blue anhydrous product was obtained. effects of lithium chloride and of ammonium thioLithium Bromide and Chloride.-The R .D.H. laboratory cyanate also were investigated in detail since pre- reagents were dried under a vacuum for two days a t 100'. liminary experiments showed that similar optical Ammonium Thiocyanate.-Hopkins and Williams Anaeffects arp produced by them. lytical Reagent (A.R.) quality ammonium thiocyanate The method used in this series of experiments is was dried under a vacuum for a day a t 100'. Hydrogen Bromide and Chloride.--The Hopkins and Wilidentical with that used in the addition of inorganic A.R. aqueous products were used after adding the acetates to cobaltous acetate in anhydrous acetic liams correct amount of acetic anhydride to react with all the acid' 4 : a set of absorption curves corresponding to water present. The acetic anhydride being added very increasing concentrations of the complexing agent slowly because of the heat produced in the reaction with is obtained. It has been pointed out by Libus, the consequent loss of the bromide or chloride. The conof bromide or chloride present waa estimated Cgniewska and Minc5 that the application of the centration by first diluting with water and adding nitric acid and excess method of continiious variations even in a modified silver nitrate, then estimating this excess with potassium ( 1 ) P 1 l'-oll 1, €T qiitrliffe and J. Walkley, J . Phys Chem , 66, 455 (1961) Rzv-nq Trans Faraday Soc., 65, 1768 (1959). ( 3 ) C E I1 n a n n R R r Kilgannon and T. K Wright, to be rnihhheti. (4) I' J Pioll znd L IT Sutcliffe, Trans Faraday SOC.,57, 1078
(1961) ( 5 ) W Lihu8, A. rgnieaska and S. Mino, Roczmki. Chemw, $4, 29 (1'360).
thiocyanate using ferric alum as an indicator. Ion Migration Experiments.-The apparatus and method used has been described previously in part I of this series of papers.' Spectrophotometry.-All equilibrium measurements were made by means of a Unicam S.P. 500 spectrophotometer fitted with a thermostated cell compartment enabling solutions to be maintained at a given temperature to within &0.05".
1'. J. PROLL ANI) L. H. STTCLIFFE
1994
1.2
i 4 0.8 3
c 0.4
Wave length, mp. Fig. 1.-The absorption spectra of Co(I1) s i t h added lithium bromide in aretir acid solution a t 26 0': A = 0.710 J I ; B = 0.533 M ; C = 0355 $1; D-= 0.17831; E = 0.0888 M ; F = 0.0144 31 lithium bromide. G is the absorption spectrum of CoRr:. The cobalt(II1 concentration is constant a t 1.82 X 10-3 M .
r--
-I
bromide in anhydrous acetic acid that the BeerLambert law is obeyed in the peak region over the concentration range studied of 1.0 X to 6.0 X 10-3 M . From the ion migration experiments it was found that although cobaltous bromide in anhydrous acetic acid contains some ionic species their actual concentration was very small, that is, the main species is the neutral molecule. Since the solution is blue in color the main complex is probably the tetrahedral CoBrz.2HOAc. This is in agreement with other workers5 who used various alcohols as solvents and assumed that the undissociated complex is of the form CoBrz.Lz where L is a solvent molecule. It therefore was concluded that the molar extinction coefficient of CoBrz.2HOAc is the same as that of cobaltous bromide in anhydrous acetic acid. This is supported by the fact that the Beer-Lambert law is obeyed over the peak region (600 to 730 mp) . The curves in Figs. 1 and 2 resulted from experiments performed on both cobaltous bromide and cobaltous acetate in anhydrous acetic acid, the cobaltous concentration being the same in both cases. For the same concentration of excess lithium bromide the molar extinction coefficient obtained was the same in both instances. At low concentrations of added lithium bromide, small differences did occur, due to the back reaction of the equilibrium Co(0Ac)p
660 700 740 K a v e length, nip. Fig. 2.--The effect of temperature on the absorption spectrum of 1.82 X iM Co(I1) TT-ith a constant lithium bromide Concentration of 0.71 111: -2 = 25'; I3 = 38.0'; C = 47.8'; D = 58.7"; E = 64.0".
620
Results and Discussion Ion Migration.-Ion migration experiments on cobaltous bromide and chloride solutions, and solutions containing excess of the corresponding lithium salt led to the conclusion that there are four likely species of the types COX+,Cox2, Coxa-, COX4'-, where X represents the bromide or chloride. The negatively charged species were favored by the addition of the corresponding lithium salt, and the positively charged species were favored by the addition of water. Spectrophotometry. (a) Addition of Bromide.Figure 1 shows the absorption curves of Co(I1) corresponding to the different concentrations of lithium bromide in anhydrous acetic acid solution. All the curves refer to a constant concentration of Co(II), namely 1.82 X loA3M . It was found from the molar extinction coefficimts of cobaltous
T'ol. 63
+ 2LiBr
CoBrz
+ 2LiOAc
This effect being more important, of course, in the situation in which lithium bromide was added to cobaltous acetate. Temperature had litt81eor no effect on the peak region of the spectrum of cobaltous bromide alone in anhydrous acetic acid principally because the dissociation constant of acetic acid is very small. Cobaltous bromide has a maximum in the absorption spectrum a t a wave length of 670 mp; added lithium bromide shifts this maximum slightly to a wave length of 696 mp and also increases the intensity of this maximum. Figure 3 shows the plot of the observed extinction coefficient against the concentration of lithium bromide a t the wave length of 695 mp. A similar plot for the acetates has been found to give a straight line,lI4but in this case a curve is obtained, which must be due to the fact that CoBr3- is a tetrahedral complex and therefore has a large extinction coefficient and is present in some quantity. The results of Libus, Ugniewska and Minc,j who used ethanol and isopropyl alcohol as solvents, have been plotted in a similar manner as shown in Fig..4. The curves are like those obtained with acetic acid as solvent. (b) Addition of Chloride.-Figure 5 shows the absorption spectra of Co(I1) with various concentrations of lithium chloride in anhydrous acetic acid solution. All the curves correspond to a constant concentration of Co(II), namely 1.82 X AI. Cobaltous chloride in acetic acid solution was found to obey Beer-Lambert's lam over the peak region as did cobaltous bromide solutions in this solvent. In conjunction with the ion migration experiments and the blue coloration of the cobaltous chloride sdutiom, it was concluded that the main species is
Xov., 1961
SPECIESOF COB.~LT(II) IT ACETIC-4CIi)
1995
CoC12.2HOhc, and that the molar extinction coefficient of cobaltous chloride in acetic acid is due mainly to this complex. The absorption maximum of cobaltous chloride in acetic acid is a t a wave length of 670 mp; addition of lithium chloride increases the intensity of absorption and shifts the maximum to 685 mp. Figures 5 and 6 are the results of experiments performed on both cobaltous acetate and cobaltous chloride in anhydrous acetic acid. For the same concentration of excess lithium chloride, the same observed molar extinction coefficient of Co(I1) was obtained for both cobalt salts. At low concentrations of lithium chloride, l4500 small deviations were noticed but as in the previous system, namely with added bromide, this was probably due to the acetate ions produced. Figure 7 shows the plot of the observed extinc400 tion coefficient a t 685 mp against the concentration I 1 of lithium chloride in acetic acid solution a t five temperatures in the range 25 to 64". Figure 8 shows the results of other workers 5,6 for different solvents plotted in a similar manner to those ob300 tained for the addition of lithium chloride and 11 hydrogen chloride in acetic acid. (c) Addition of Thiocyanate.-Figure 9 shows the absorption curves of Co(I1) corresponding to 0.2 0.2 0.6 0.8 various concentrations of ammonium thiocyanate in Concn. of lithium bromide, M . anhydrous acetic acid solution. All the spectra Fig. 3.--The dependence of the molar extinction coefwere measured with a constant concentration of ficient E a t 695 mp on the concentration of lithium bromide Co(II), namely 1.82 x M . Cobaltous thio- a t the temperatures: -4 = 25.0"; B = 38.0'; C = 47.8"; cyanate in acetic acid solution obeys Beer-Lam- D = 58.7'; E = 64.0". bert's law over the peak region; the solutions are blue in acetic acid, hence it was concluded that the main species is the neutral tetrahedral molecule, 800 i namely, C O ( C N S ) ~ . ~ H O Athere C , being two solvent molecules as in the previously discussed bromide and chloride salts of cobalt(I1) in acetic acid -. 600 solution. The spectrum of cobaltous thiocyanate E' in acetic acid (see Fig. 9), shows one absorption band in the visible region, the wave length of the 400 maximum being 620 mp. The addition of ammonium thiocyanate increases the intensity of the d+d u; transition but there is no shift in the wave length 200 of this maximum. Figure 10 shows the plot of the observed extinction coefficient a t the wave length of 620 mp against the concentration of ammonium l-LpL--_-L _i_...-__ thiocyanate a t a temperature of 25.0'. 0.5 1 .o 1.5 2.0 The effects of fluoride, cyanide and azide also were Concn. of lithium bromide, 31'. tried. The fluoride used was the ammonium salt 4.-Plots of the observed extinction coefficient 6 but this is not very soluble, and did not show any at Fig. 695 mp against the concentration of lithium bromide. visible effect on the spectrum of cobaltous acetate. ( A ) our results using acetic acid as solvent a t 25.0'. (B) Potassium cyanide was found to be very soluble and ( C ) the results taken from the work of Libus, Ugniewska but its addition to cobaltous acetate in anhydrous and Mints using (B) isopropyl alcoho; and (C) ethanol as a t a temperature of 19 f 1 . 0 are the experiacetic acid led to some precipitation and hence the solvents mental data and X are the calculated results. system could not be studied by spectrophotometry. Azide was found to affect the spectrum of cobaltous acetate in anhydrous acetic acid in that the peak tion constant of HNZ in acetic acid is very much was shifted to longer wave lengths and there was an less than that of acetic acid itself. increase in the intensity of absorption. It was Discussion found, however, that the effect was the same as Calculated curves similar to the experimental that of sodium acetate, hence it was concluded that curves shown in Figs. 3, 4, 7, 8 and 10 may be genthe equilihium erated by considering the effect of the anion X- on S a N 3 + HOAc NaOAc HNs the corresponding cobaltous salt COXZas follows. is present. From this it appears that the dissocia- For simplicity, molecules of solvation are omitted and all the cobalt species are assumed to be of the (6) W.D. 13eaver, L. E. Trevorrow. W. E. Estill, Pi C. Yatea and tetrahedral configuration. T,E. Maare, J , Am, Cham8 SOC.,76, 4556 (10531.
I1
1
1-
a
+
--
~
P. J. PROLL AN.) L. H. SUTCLIFFE
1996
1-01.65
500
400
.. I
-
300
c)
I 4 ui
200
0.2
I
I I 610 650 Wave length, mp.
I
I
690
Fig. 5.-The absorption spectra of Co(I1) with added 0.2 0.4 0.6 lithium chloride in acetic acid solution a t 25.0", the conConcn. of lithium chloride, M . centration of Co(1I) being 1.82 X lo-* M . Concentrations Fig. 7.-The plot of the observed molar extinction coefof lithium chloride are: A = 0.500 M ; B = 0.250 M ; C = 0.125 M ; D = 0.0625 M ; E = 0.0313 M ; F = 1.82 X ficient E a t 685 mp against the concentration of lithium chloride a t the temperatures: -4 = 25.0"; I3 = 38.0'; M colmltous chloride. C = 47.8": D = 58.7": E = 64.0'.
600
610 650 600 Wave length, mp. Fig. 6.-The effect of temperature on the absorption spectrum of co(I1) of concentration 1.82 x 10-8 M , with x constant concentration of lithium chloride of 0.500 64. A = 25.0 ; R = 38.0"; C = 47.8'; D = 58.7'; E = 64.0'.
+ x- K2 Ka cox3- + xcox2
cox,-
*
(1 )
h
+ *KeIX-l + cK2K:[X-12 + KzLX-1 4- &&[X-l2
1
I
---
'
I
0.2 0.4 0.6 0.8 Concn. of monovalent chloride, M . Fig. 8.-The observed molar extinction coefficient E a t 685 mp us. the concentration of ( A ) hydrogen chloride and (B) lithium chloride in anhydrous acetic acid a t 25.0"; (C) monobutylammonium chloride in isopropyl alcohols a t 19 f 1" and ( D ) lithium chloride in 2-octnnol~a t 30".
where c2, € 3 and e4 are the molar extinction coefficients of the solvated species COXZ,COX$- and COx4'-, respectively, a t a given m v e length. Since the dissociation constant K x of salts in anhydrous acetic acid is known to be T7er.y small7 we may replace [X-] by ( K x[l\rX])l *,hence e q i ation 3 becomes
cox4*-
from which it may be shown that the observed extinction coefficient E at a fixed wave length is given by E =
r
(3)
For all the additions reported in this paper was identified with the observed extinction coefficient for solutions of Cox2 with no added halide salts. (7) S. Bruckenstein and I. hf. KnlthotT, J . Am. C h m . Soc., 78, 2974 (1956).
SPECIES OF COB.ILT(II) IS ACETICB C ~ D
Yov., 1961
1997
Consicieriiig first the bromide bybtein, ai1 empirical expression, namely €
+ 4600[ LiBr] + 900 [LiBr] + 5.O[LiBr]'/z + l.O[LiBr]
200 1.0
= --___-
'12
(5)
was found t o gim a curve which shows reasonable agreement with the experimental curve a t 25.0" (see Fig. 4A), from which it was deduced that the values of e3 and E ( at 695 mp are 920 and 900 J1-l cm.-l, respectively. The dissociation constant of lithium bromide i i i acetic acid ( K B )has been estimated to be 7.2 x J I at 30°,8but this value is unreliable since it waq found from E'uoss-Kraus plots of con4 ductivity data (see Bruckenstein and Kolthoff) .' Assuming a posit:ve enthalpy change for K B , a lower value than the above is required a t 25'. A 0.2 value a t 2.5" of 5 X low7111 was used for K B and hence we were able to evaluate K1 and KB to be 7 X lo3 and 3 X lo2 M-l, respectively. The reliability of the values will, of course, be subject to some doubt until a more dependable value of K B is determined. The values for the equilibrium conI 1 1 stants K z and K 3 are not unreasonable, since in the 600 650 700 case of the addition of inorganic acetates to cobalTTTave length, mp. tous acetate in anhydrous acetic acid1g4the product Fig. 9.-The absorption spectrum of Co(I1) with added K2K3for the equilibria analogous to (1) and (21 ammonium thiocyanate in acetic acid solution a t 25.0". above has a value of 2 X lo351-*,being composed Concentrations of ammonium thiocyanate are: A = M ; B = 0.389 M ; C = 0.265 &I; D = 0.133 M ; of K z = :!0 J - 1 and Zi, = I x l o 4 A-l. The 0.530 = 0.0663 M ; F = 0.0076 A/ with the concentration of Co values of e3 and are probably of the right order E (11)being constant a t 1.82 X 10-3 hf; G = 1.82 X ill since a t tbis wave length the extinction coefficients cobaltous thiocyanate. of cobalt(I1) with a very large excess of lithium I 2000 bromide ill ethanol or isopropyl alcohol are lo~ver.~ I Treating the results for alcohol solutions5 in a similar manner leads to the conclusion that the value of €4 is 900 M-' cm.-' in both of these sol1600 vents, as it is in acetic acid a t the same wave length of 695 nip. Since this is the value of the extinction coefficient of CoBr42-, it mould not be expected to change with solvent as no solvent ligand 7 1200 is involved, as it is tetrahedral. In the case of CoBr3-, e 3 was calculated to have values 775 and 7170 J1-l mi. respectively, for isopropyl alcohol and ethanol compared with 920 111-l em.-' for acetic acid. This is not unexpected since it is the extinction coeflicient of the CoBr3- ion, which must have one c:olvent ligand attached to it for the tetrahedral configuration. It also was concluded that 2 1.3 with ethawl :is solvent the value of K z K ~ ' / is Jf -'I2 and of K 2 K J Kis~ 1.0 JI-1; with isopropyl alcohol as solveqt the corresponding values arc 3.0 i i 1 - ' / 2 and 1.0 U-l, respectively, in the temperature range 18-20°. The agreement between the calculated and observed points is shown in Figs. 4B 0.2 0.4 0.6 and 42. The values of K B in these solvents are C'oncn. of ammonium thiocyanate, AI. unknown, hciice it is not possible to evaluate K 2or Fig. 10.-Thc plot of the ohserveti cxtiIirtion cw#icient Zi, for either system. The value of the observed at 620 nip ag:tinst the coriccritration of :mmonium thiomolar extinction coefficient of cobaltous bromide cyanate a t the temperature 25.0". o are the obscrvctl in each of the three solvents mentioned previously cxpcriiiientd points and X arc the colculnted points. a t 695 mfi is much greater than that of cobaltous KZKB'" and Z S a > K. The only measured value for the dissociation constant of hydrogen bromide in acetic acid is that found by Smith and Elliot9 which is known to be inaccurate owing to the use of Fuoss-Kraus plots. The value found was 2 X lo-' M at 25") which would not fit in a regular order with the results of other workers8 for sodium, lithium and potassium, even allowing for the variation of temperature. I t might be argued that the forward reaction of the equilibrium LiBr
Vol. 65
+ HBr
KzKIKc(M-') 2.5 1.5 1.2 0 8 0.7
6 0 5.5 5.0 4.8
Because the expression for the observed extinction coefficient with the addition of chloride is very dependent upon the dissociation constant of the chloride concerned it was decided to try the addition of another chloride. As in the case of the bromides, the only other monovalent chloride that is very soluble is hydrogen chloride. The addition of hydrogen chloride in anhydrous acetic acid to cobaltous solutions in the same solvent had the same effect as the addition of lithium chloride in that the shape of the spectra obtained were identical, but the observed molar extinction coefficients for the same concentrations of both chlorides were lower for hydrogen chloride than for lithium chloride. This is shown in Fig. 8, together with the results obtained for alcohol solutions by other workers5>6when replotted. By choosing suitable values for the constants, the plot of the observed extinction coefficient against the concentration of HC1 can be fitted and hence the ratio K C / K H Ccan be found, where K K Cis the dissociation constant of hydrogen chloride in anhydrous acetic acid. The value of this ratio is 12. Kolthoff and Bruckenstein' using potentiometric methods have determined Kc: using KHc determined spectrophotometrically*Othe value of the ratio K c l K ~ cis 26. The agreement is reasonable when one considers the difficulties involved in such work. For example, the effect of water in our experiments would lead to an apparent decrease in the value of the dissociation constant concerned, whereas in the potentiometric methods an apparent increase mould be observed. For the system cobaltous chloride plus monobutylammonium chloride in isopropyl alcohol5 the following empirical relationship was found
is of importance since the effects of HBr and LiBr are so very similar in magnitude, but it can be 105 + 5600[C4HgNH3Cl]'1% 6500[C4HgXHaCl] € = shown that the equilibrium constant for this is 1.0 + lO[C,HgNH&I]'/2 lO[C4HsNH&IT equal to K B K A ~ I K A K Hwhich B is of the order of This leads to the conclusion that the values of E:$ 10-8, where K A and K A are ~ the dissociation con- and 6 %are 560 and 650 M-l cm. -l, respectively, and stants of lithium acetate and acetic acid, respec- the values of K 2 K ~ ' / and 2 K a 2 K (where ~ K N is tively, and the other symbols are as previously the dissociation constant of C4H9NH3C1in isopropyl designal ed. alcohol) both have a value of 10. Since K N has not Considering now the chloride system by choosing been determined only the ratio K2/K3 can be obsuitablr values for the constants, calculated plots tained; the value of this is 10. I n the case of the again can be made to fit the experimental curves. addition of lithium chloride to cobaltous chloride By finding the variation with temperature of Kz in 2-octano16an empirical equation, namely KC'/^ and KzK3Kc (see Table 11) it is possible to 63 + 4480[LiC1]'/2 65O[LiC1] estimate the over-all enthalpy changes for each of E = 1.0 + 8.01LiC1l1/?+ l.OlIm these combined processes. The values found were was found t o fit the experimental data, from which -2.1 f 0.5 and -6.2 + 2 kcal. mole-l. The value of K c , the dissociation constant of it was calciilatcd that the values of e3 and €4 are 560 lithium chloride in acetic acid a t 23.0", has been and 650 J1-l cm.-', respectively, and the values of determined,' hence it is possible to evaluate both K&O1/2 and K3K2Ko are 8.0 M - - " 2 and 1.0 J - l , Kz and K 3 at this temperature. The values found respectively, where K Ois the dissociation constant were 2.5 X lo4 and 1.2 X lo3 M-l, respectively. of lithium chloride in 2-octanol. As for the bromide system the value of €4 does c3 and e4 were found to have values of 600 and 650 not appear to change from solvent to solvent, but M-I cm.-l, respectively, a t 685 mp.
+ +
+
(9) T. L. Smith and J. H. Elliot, J . Am. Chem. Soc., 7 6 , 3566 (1953).
( I O ) I 11 Kolthoff and S. Bruckensteln. J A m Chem. Soc., 78, 1 (1956)
Kov., 1961
SPECIESOF COBALT(II) ISACETICACID
1999
once again the value of e3 does change. In the al- ably tetrahedral for the acetate, also has been cohol solvents a value of 560 M'--l ern.-' is obtained shown to be formed. These anionic species of coand a value of 600 11f-l cm.-l in acetic acid, which balt(I1) are both related to the neutral species by is due to one solvent molecule being included in equilibria (1) and ( 2 ) . this species. The order of magnitude of the terms The neutral species for the acetate i b octahedral, KzKp~'/z and K2K3K~,and K2K01/2 and K2K3K0 while for the other ligands it is tetrahedral. The are similar to those obtained with acetic acid as sol- spectra observcd for all the species arc very similar vent. in that they show only one main absorption band For the thiocyanate system a calculated plot in the visible region of the spectrum which is a d + fitting the experimental curve was obtained (Fig. d transition. For the neutral species the wave 10) with the values of t3 and e, of 1900 and 2500 lengths of the absorption maxima fall in the order cm.-l, respectively, and KZKT"~ and K 2 K 3 K ~ bromide (672 mp), chloride (663 my), thiocyanate of 23 J-l'z and 1.0 X-', respectively, at the tem- (620 mp) and acetate (526 mp). In x-iew of the perature of 25.0'. The individual values of the broad nature of the band and its large extinction constants K2 and K 3 cannot be determined since coefficient, the transitions must be spin allowed and the value of KT, the dissociation constant of am- must be the + 4E'1 or the 4 A 2 y transition in monium t hiocyanate in anhydrous acetic acid, is the case of the latter anion since it is of octahedral not known, but the ratio K2/K3 mas found to be configuration, and probably the 4F2c 4 X 2 transition 530. The value of KT is not expected to be very in the other cases since they are of the tetrahedral different from the bromides and chlorides and configuration. The addition of excebs of acetate, should be in the range to lo-$ M and hence the thiocyanate, chloride and bromide shifts the maxiorder of magnitude of the product K2K3 will be the mum of the absorption spectrum to longer xave same as that found for the bromide and chloride lengths in all these systems, the maxima being a t systems, that is, lo6 to lo9 36.5, 620, 685 and 69.5 m p , respectively. To a first I n the experiments on the addition of acetates, approximation thib can be assumed to he the order described in part I, the effect of the addition of of the decreasing Dq d u e s for Co(I1) in a tetrahewater was found to be attributable to a dielectric dral field since the transition is the same for all these constant change on the equilibrium ligands. The extinction coefficients obtained for the tetrahedral anionic species of Co(I1) a t the K3 above maxima are shown in Table 111. ("o(0iic)g- + OilcCO(OAC)~~In the present system, the addition of water was found to reduce the molar extinction coefficient of Co(1I) in the presence of excess lithium salts, but this is more likely to be due to the increased ionization of the acetic acid, since the blue color is completely destroyed with relatively small amounts of water. The addition of sodium acetate produced the same effect. The significance of the results in terms of the catalytic activity now can be discussed. The data given in references 2 and 3 shows that reactivity is a t a maximum when the ratio of Br-:Co(II) is either 1: 1 or 2 :1, depending upon the experimental conditions. From the results obtained in this paper, th(. cobaltous species would be expected to be of tetrahedral structure and can be tentatively given the formulas (CoBr.3HOAc)+ and CoBrz. 2HOAc. The dissociation of the COB^^^- complex on increasing the temperature should favor the production of the latter complex thereby increasing the catalytic activity of the system. The poorer reactivity of the ion CoBr42- might be due to the difficulty of producing Co(III), which is not likely to exist as a tetrahedral complex. Tetrahedral complexes with fewer bromine ligands might be more amenable to group transfer of say an acetoxy radical to give octahedral Co(II1). General Conclusions The addition of acetate, chloride, bromide and thiocyanate, to cobalt (11) in acetic acid solution has been shown to lead to the formation of tetrahedral species of the type COX^^-, where X represents the ligand$ mentioned above. A species of the type COXg-, which is tetrahedral for the chloride, bromide and thiocyanate complexes and prob-
TABLEI11 Extinction coefficients (?.I-1 cm -1) of CoXa- CoX4'-
Additlie
Acetate Bromide Chloride Thiocyanate
920 600 1900
1700 900 650 2500
Estd oscillator strength of CoX42-
Wave Hdlflength of band max. uidth. em.? L ~ -1 I
17700 11.200 14600 16130
1150 32 5 X 1810 i.5 X 22i0 6 85 X l o e 3 1950 22 5 X
To a first approximation the Dq values should be in the same order as the oscillator strengths which can be calculated from the extinction coefficient and the half-band Jvidth. The oscillator strength F being given by F
=
4.32 X
lO-9Jt
d, = 4.6 X 10-9
€0
dw
where eo is the maximum extinction coefficient and dw is the half width in cm.-' (reference 11). Since all the additives cause the formation of both Cox3and the resultant spectrum obtained is a mixture of these and that of the neutral molecule, but the absorption spectra do not change appreciably in shape after the addition of a small excess of the additive. Hence to a first approximation the oscillator strengths of the COX^^- can be calculated using the half band width as estimated from the spectra obtained at very high concentrations of the additivcs. The results of these calculations are shown in Table 111. The calculated oscillator strengths fall in the order C0(0Ac)4~-> C O ( C S S ) ~> ~ -CoBr2- > CoC O C ~ ~ ~The - . order of decreasing Dq values found from the order of the wave lengths of the maxima in the visible region is as above with the bromide and chloride interchanged. The fact that the oscillator (11) C. K. J$rgenben, Acta Chem Scend., 9 , PO7 (1955).
E. BLOMGREN, J. O'M..BOCKHIS AND C. JESCH
2000
T'ol. 63
TABLE 1V Additive
Acetate Bromide Chloride Thiocyanate
Values of KtKxIl*and KzKsKx a t 25.0'
K ~ K ~ '=/ Zo 01 ~ 4 - ' / 2 K ~ K ~ K= A0.049 M - ' K2KB'/:= 5 . 0 M-'/a K ~ K ~ K= E1 .O M-' K ~ K c ' /=~ 7 . 0 M-'12 K z K ~ K c 2 5 M-' K~KT'= ~ ' 2 3 . 0 M-'" K ~ K ~ K=T1 . O M-'
Over-all enthalpy change, kcal. mole-'
.......
-
-12.0 f 4 . 0 - 2 1 lt 0 . 5 - 6.2 f 2.0
........ *.
strength for CoBr4*- is too large is probably due to the absorption spectrum of cobalt(I1) in the presence of excess bromide being much less Gaussian in shape than for the other systems. The value of the oscillator strength of the CoCld2- species calculated above of 6.85 X agrees reasonably well with the value of 6.0 X quoted by Balehausen and Liehr,lZparticularly in view of the approximations involved in these calculations. The oscillator strengths when plotted against the position of the absorption maximum (cm. -') are reasonably linear, as one would predict. The good agreement helps to confirm our interpretation of the spectral effects due to the addition of compounds containing the various ligands. Table I V shows the over-all enthalpy changes, the over-all free energy and over-all entropy changes a t 25.0' for the combined processes KzKx'/' and K 2 K 3 K x , where K x represents the dissociation coii(12) C . J Balehausen and B . D. Liehr, (1958).
J. M o l .
1.8 i 0.1 2.2 i0.5
Spectroscopy, 2 , 3 4 2
Over-all free energy change a t 25.0'. kcal. mole -1
Over-all entropy change a t 25.0°, cal. mole-' deg. - 1
......... 1.8 f 0 02 -0.96 f .04 0.00 i .04 -1.16z.k .04 - 0 . 5 5 lt .04 -1.86 i .04 0 00 lt .04
........ 0.0 f 0 . 4 - 4.0 f 1.7 -40 i 14 - 3.1 i 1.7 -19 lt7
........
........
stant of the inorganic additives, where KA is the dissociation constant of sodium acetate, K B and Kc are the dissociation constants of lithium bromide and chloride, respectively, and KT is the dissociation of ammonium thiocyanate, in all cases acetic acid being the solvent. Since it is reasonably safe to assume that increase in temperature will increase the dissociation of the inorganic additives, the enthalpy change for this process is positive. With this assumption it is possible to estimate the enthalpy change for Kz for both the bromide and chloride systems to be not greater than -2 kcal. mole-l, and that for Ky to be not greater than -10 =t 4 and -4 f 2 kcal. mole-1, respectively. For the bromide and chloride systems it also was found possible to evaluate the difference in the enthalpy changes ( A H z - A H 3 ) for the processes KZand K3. This was found to be for the bromide $7 =t 3 kcal. mole-' and for the chloride 3.3 i 1 kcal. mole-'. The authors are indebted to Dr. J. II. Binks of this Department for much helpful discussion.
THE ADSORPTIOS OF BUTYL, PHENYL AND NAPHTHYL COMPOUSDS A T THE IKTERFACE MERCURY-AQUEOUS ACID SOLUTION BY E. BLOMGREN, J. O'M. BOCKRIS AND C. JESCH John Harrison Laboratory of Chemistry, Cniversity of Pennsylvania, Philadelphia 4, Pa. Received April 87, 1961
The adsorption on mercury from acid solutions of butyl, phenyl and naphthyl compounds having the groups OH, CHO, COOH, CN, SH, S, CO, yH3+ and SO3- has been examined as a function of concentration and potential. The electroca illary method was applied. The following quantities have been calculated thermodynamically: the net charge on the soktion; the degree of coverage of the surface with organic compound, the net free energy of adsorption of organic adsorbate as a function of coverage and potential; the adsorption of H30 and C1- ions in relation to t h r adsorption of organic molecules. The effect of diffusion on the rate of attainment of adsorption equilibrium is analyzed. Naphthyl compounds have a 103- 10' times higher adsorbability than corresponding butyl compounds, with phenyl compounds intermediate. For substances with the Rame hydrocarbon radical, the adsorbability varies with the substituent group by a factor of up to lo4. Aromatic cations show a special behavior. The orientation of the adsorbed molecules has been determined, and t h r character of the adsorptive forces is discussed in terms of metal-adsorbate and solution-adsorbate interaction. A method is proposed bv which the intrinsic free energy of adsorption of a hydrocarbon radical in an organic substance upoii a given metal can be derived. I t is about 6 kcal./mole greater for the naphthyl than for the butyl radical The degree of effectiveness in inhibiting the dissolution of iron and the degree of coverage on mercury show a parnllelisin for the majority of compound5 studied. +
GOUY'S' examination of the surface tension depression of mercury by organic compounds does not permit calculation of surface coverages. Frumkin2 determined surface tension depressions for (1) G. Gouy, Ann. chirn. p h y s , [7] 29, 145 (1903): [SI 8, 291 (1906); [8], 9, 75 (1906); Ann. phys., 191 6, 5 (1916); [91 1, 129 (1917). (2) A. N. Frumkin, Z. Physik. 36, 792 (1926).
various concentrations of t-amyl alcohol in aqueous solutions and interpreted the results in terms of a semi-empirical adsorption isotherm. Conway, Bockris and Lovrecek3 used the electrocapillary method to examine the adsorption of alkaloids on (3) B. E. Conway, J. O'h'f. Bockris and B. Lovrecek. C I T C E Proc., 8, 207 (1955).
