Jan., 1952
THECOPPER(II)-CYANIDE REACTION
ethylenetetramine complex ion, is indicative of the strain present in the ion. The determination of the formation constants a t two temperatures permits calculation of the free energy of the ions and approximate heats of formation of the complex ions. These are given in Table IV. A comparison of the data of the first and second step equilibria in the complex ions containing triethylenetetramine and diethylenetriamine shows that the increase in stability in the second step equilibrium is very much smaller for the trien complexes than for the dien complex ions. This is to be expected since the second step of the trien complexes involves formation of dinuclear complex ions. No evidences of the hydrogen complexes reported by Schwarzenbach were obtained in this investigation. Irving and Williams‘o have recently reviewed the stability of complex ions of bivalent metals. For the ions, Cu(II), Ni(II), Co(II), Zn(I1) and Cd(I1) the stability decreases in that order irrespective of the nature of the coordinating group. (10) H.Irving and J. P Williams, Nature, 162, 746 (1948).
19
Calvin and Melchior” obtained the same order of stability. Attempts by these authors to relate the order of stability to some property of the metal atom or ion showed that the correlation was best obtained between the second ionization potential of the gaseous atoms and the relative stability of the complex. Figure 3 shows the logarithms of the stability constants of complex ions with various amines with a coordination number of four which have been studied quantitatively along with the second ionization potentials as given by Latimer.12 Since only a very few data are available for Fe(I1) and Mn(I1) complexes with this coordination number these values are not included but the decrease in the second ionization potential observed for these ions is also found in the log IC data of their complexes. Similar results are indicated by the meager data available for complexes with a coordination number of six. (11) M. Calvin and N. C . Melchior. J . A m . Chum. Boc., 70, 3270 (1948). (12) W. M. Latimer, “The Oxidation States of the Elements and Their Potentisla in Aqueous Solutione,” Prentice-Hall, Inc., New York, N. Y., 1938. pp. 14-15.
COMPLEXES IN OXIDATION-REDUCTION REACTIONS. THE COPPER(I1)-CYANIDE REACTION BY FREDERICK R. DUKEAND WELBYG . COURTNEY Contribution No. 1.99from the Institute for Atomic Research and Department of Chemistry, Iowa Slate College, Ames, Iowa Received August 30, 1061
Many homogeneous oxidation-reduction reactions have been shown to involve coordination com lexes as intermediates. In the present work, copper(I1) is shown to coordinate four cyanide ions, this complex ion then yieling CU(CN)~”and CN radical; high concentrations of ammonia are present in the reaction mixture in order to compete with the cyanide ions for the coordinat,ion positions on the copper(11)ion, thereby bringing the rate of the reaction into a measurable range, A possible reason for the necessity of four cyanide ions in the reacting complex is discussed.
Out of the studies on the mechanism of electron transfer have come some promising possible generalizations. One of these might be stated as follows: coordination complexes are involved mechanistically as intermediates in homogeneous ionic oxidation-reduction reactions. This hypothesis appears to be particularly applicable to reactions involving cationic oxidants and anionic or “Lewisbase” reductants. For example, systems in which such complexes have been observed as intermediates are Fe(111)-I -, Mn (111)-C2Or, 2, Ce (1V)glyc01,~ Ce(IV)-CI-,5 Fe(III)S03-,6 Ce(1V)CzOa,7 and a number of others. A point of further interest in connection with these coordination intermediates is any relationship which might exist between the number of oxidizable anions in the complex and the rate of “electron transfer.” For instance, the complex Fe12+yields ferrous ion very much more rapidly than does the (1) A. V. Hershey and W. C . Bray, J . A m . Chem. Soc., 08, 1760 (1936). (2) F. R. Duke, ibid., 69, 2885 (1947). (3) H.Taube, ibid., 70, 1216 (1948). (4) F. R. Duke and A. A. Forist, ibid., 71, 2790 (1949). ( 5 ) F. R. Duke and J. Anderegg, unpublished data. (6) F. R. Duke and A. Bottoms, unpublished data. (7) S. D. Ross and C. G . Swain, J . A m . Ckem. Soc., 69, 1325 (1947).
complex FeIf2; Mn(C20a)+ yields manganous very much more rapidly than does Mn(CzO&-. The present work examines the Cu(I1)-CNreaction from the points of view outlined above.
Experimental Reagent grade chemicals were used. Acidified cupric chloride stock solution was standardized through thiosulfate against KpCraOTwith a small quantity of Na2COa added, and potassium cyanide was standardized against AgNOI with the iodide end-point. The ammonia was cooled, diluted, standardized against HCI, and thereafter kept in an ice-box, as were all solutions containing ammonia. Kinetic runs were made at 0’ f 0.lo, maintained by a bath of melting ice in a dewar flask. Scparate erlenmeyer flasks containing 50 ml. of the desired cupric ammonia solution and a cyanide solution of appropriate concentration were cooled to equilibrium in the icebath. Five ml. of the cyanide solution were then added to the rapidly shaken copper solution by means of a pipet calibrated to give 5.00 f 0.02 ml. Five-ml. portions of the reacting solution were removed by a calibrated pipet at known times and quenched in 2 ml. of a solution containing 0.2 M zinc nitrate in approximately 10 M ammonia. The optical density a t 600 m r of the quenched solution was observed and compared against an empirical curve constructed from the optical densities of known cupric copper concentrations undcr identical conditions except for the colorlcss cuprous cyanide complexes. The concentration of total cupric copper in the reacting solution was thus determined. Concen-
FREDERICK R. DUKEAND WE~BYG. COURTNEY
20
trations of other ions subsequently cited likewise refer to the reacting solution. The ionic strength of the reacting solution was maintained at 1.0 with sodium nitrate. The ammonia concentrations remained sutliciently standard over 18 hours. and the cyanide solutions for even longer. Optical measurements were performed on a Coleman Universal Spectrophotometer, Model 14. A separate study was made to determine -+e cause of the distinct violet color resulting from the addition of c y m d e to ammoniated cupric copper. The optical density at 014 mp of the copper-ammonia-cyanide rmxture was measured as a function of time, since it is unstable. Extrapolation over 20 seconds to zero time gave 4, the extinction at zero time. The above wave length was chosen from fast test rum as the approximate maximum in the absorption spectrum of the violet smcies. These ootical measurements were made with a Gary Recording-Spectrophotometer, Model 12, Serial 15.
Results and Discussion It was noticed that a shift from blue to blueviolet occurred when CN- was added to ammoniacal cupric solutions. The violet color absorbed most strongly light of wave length 615 mp. The effect of varying CN- a t constant Cu++ on the extinction was determined a t this wave length and found to be lhear in CN-. Since K = Cu(CN);-'/(Cu++) (CN)' and CU(CN),~-X= KE, where E is the extinction, K(Cu++)(CN-)x = KE, or at constant Cu++, (CN-)X = K'E. The straight line variatian of E with (CN-) indicates that x = 1and that the inconsequential amounts of copper and cyanide are stored aa CuCN+ a t high ammonia concentrations. Thus we were able to conclude that the Cu++ could be represented by total copper concentration and that CN- initially was equal to total cyanide concentration. The stoichiometry of the reaction was found to be 8CN2Cu++ 3 ~ C U ( C N ) ~ =(CN)z. Cyanogen was found to hydrolyze slowly compared with the rate of the oxidation reaction. Plots were prepared of Cu++ us. time. Some typical curves are shown in Fig. 1. From these
+
+
-ai
I
-5
P -6
-70
.I
Fig. 2.-Plot
2
4 5 log a.
6
IT
.8
9
ID
for the determination of kinetic order in copper (11)and cyanide.
time ratio of CN- to Cu++ is a (Fig. 2). These plots were found to have a slope of 4 for a variety of values of a, and intercepts appropriate t o the equation: -R, = kTZI1a4. This result indicates a rate expression - Ro = - (dThII/dt)o = kll(Cu++)o(CN-)'. (1)
Next, the effect of variation in (NEb) was considered. Utilizing the cupric-ammonia stability constants KA where KL = [Cu(NEI3)~'I/[Cu(NH3)i1]m:a], and the relation T C ~ I I= [Cu(")I:' i[Cu(NEb):'], since the concentrations of less-ammoniated species can be calculated to be insignificant and that of the cyanide complex is assumed to be insigdicant, it can easily be shown that
+
Do12
-R I
500
I
1000 SECONDS.
1500
I
2000
'
in total copper (11) with time at various .cyanide concentrations.
plots, -d Cu/dt was determined by the plane surface mirror technique, both a t zero time and other times. The zero time rates are uncomplicated by the presence of cuprous copper and were therefore analyzed first. Plots were made of log (- R,) urnsus log a for constant original copper(II), T&II, where R, is the initial reaction rate and the zero
* *
--
-
(3)
- ~T,II)'
(4)
Equation (4) reduces t o (1) at eero time. Integration of (4) gives 1
, \
-
kb:a, - k: -k;&, kR' ( T C u d ( T & - 4T&1
k:
Fig. I.-Variation
.3
where (Id", is the activity of the ammonia. Also, since (CN-) = TCN-= TEN- - 4Tcu11 ~T&II from stoichiometry, the integrable rate equation is - E = kRII(ThII)(T&- - 4T&XI f 4Tcu11)'
.0013
0 0
Vol. 56
+
- 242%+ 64zs - (In 2,1 f 1220 242: f 64%;) = kg'(T&- - 4y&1)4 x t
In ; 122
+
(5)
where x = Tcu1i/T&- - 4T&1 4TcUI1)and G refers to zero time. Equation (6), X - X , = K't, corresponding to (5) is valid for all conditions other than TgN- - 4T&11 = 0, when (4) must be reintegrated . The slow reaction may involve Cu(CN)4@&)rather than Cu(CN)d- as the disproportionating species, and in this case (2) is replaced by
Jan., 1952
THECOPPER(II)-CYANIDE REACTION
sent activated complexes of considerably higher free energy than that involving four cyanide ions.
K' now becomes
14
A similar treatment could be given for an intermediate complex of Cu(CN)c(NHa)a-. The validity of the mechanism which has been postulated may be experimentally verified by the use of (6), for the product ICK' should be independent of copper(I1) and cyanide concentrations. The role of ammonia in the intermediate complex is carried m K', per equation (4) or (8). For the excessive ammonia concentrations, K' is a true constant;->Plots were prepared where the left side of equation (6) was calculated and graphed us. t. Figure 3 exemplifies these plots. The lines which resulted were straight and had the slopes, K", indicated in column 4 in Table I. TABLE I
PSEUDO AND TRUH CONSTANTS POR p = 1, T = 0" 12.6
11.2 10.5
0.00125 0.0150 .00125 .0100 .00150 .0120 .00150 .00900 .00200 .00800 .00250 .0100 .00300 . m o o .00600 .0120 .00600 .00900 .00600 .00600 .00300 . m o o .00300 .00600 .00200 .00600
21
5.3 X 4.8 X 4.9 X 7.0 X
106 106 106
5.3 X 4.8 X 6.3 X 5.1 X 9.9 X 9.2 X 1.2 X
10' 106 106 10s 10' 106 106
106
....... .......
2.8 X 10" 2.7 2.7 3.6 3.0 2.9 2.8 2.5 3.3 2.7 2.8 X 1025 2.8 2.9 X 102'
3.3 x low 3.2 3.2 4.2 3.6
3.4 3.3 3.0 3.9 3.2 4.0 x 1020 4.0 4.1 x 10x8
The last two columns of Table I indicate the results of an attempt to determine the role of ammonia in the reaction by calculating the denominator of K' in the two mechanisms mentioned previously. If we now substitute the formal concentrations of ammonia, TNR:,for activities, the denominator of K' in (4) becomes 2.4 X los T i H I 1.5 X 10l5 TRHIor in (8), 3.6 X 1Olo TkRI 2.3 X 1Olo TRHI. The values of ICKII and ~ K I in V Table I are calculated under these conditions and should therefore be considered as independent of variations in copper(II), cyanide or ammonia. These values definitely disallow any categorical statements regarding the function of ammonia in the reaction, except that ammonia greatly lowers the rate in closest agreement with the assumption that Cu(CN)*- is the intermediate species. The high inverse dependence of the rate on ammonia concentration precluded the possibility of obtaining data over wide ammonia variations. In searching for a reason why four CN- ions must be in the complex for the reaction t o proceed, we tentatively arrived a t the conclusion that cuprous cyanide complexes having less than three cyanides must have very high free energy, too high to represent to any measurable extent the kinetic product of the reaction. Another way of saying essentially the same thing is that the activated complex must have substantially the properties of the product, and the high energy products repre-
+
+
12
10
$ 8
I H
6
4
2
0 0
500 Fig. &-Plot
lo00 1500 2000 Seconds. testing equation (6).
2500
This work was performed in the Ames Laboratory of the Atomic Energy Commission.
REMARKS ROBERT E. CONNICK: The authors have stressed an interpretation of the data in terms of an activated complex structurally very similar to known or plausible complex ions. I t should be pointed out that such a picture is not unique and the activated complex may consist of an arrangement of atoms uite unlike the stable complex ion. The kinetic data will be%tted equally well as long as the stoichiometric composition of the activated complex is the same. Thus one could interpret a first order rate law for the decomposition of the ceric-glycol complex in two steps: first, the dissociation of the complex into ceric ion and glycol; and, second, the collision of these two species to form the activated complex. The geometrical arrangement in space of the ceric ion and glycol might be completely unlike that of the stable complex. Such an interpretation, when applied to the co per cyanide reaction gives a plausible rationalization of t i e observed fourth order dependence on cyanide ion concentration. Thus the activated complex may consist of the copper with three cyanide groups arranged around it in positions similar to those found in the product, Cu(CN)r--. The fourth cyanide ion would be the one oxidized and could have a position completely unrelated to any stable complex ion. For example, it may be turned 180"from its normal coordination configuration. It is of course imDossible to deduce from kinetic data the structural arrangement of the atoms within the activated complex. REPLY:My attitude on this uestion is consistent with the principle of accepting the sim&r of two possible alternatives which cannot be distinguished. Further, the existence of an interatomic orbital involving the oxidant and reductant could rovide a convenient path for the electrons. Perhaps compehng evidence will show up in the future in favor of one or the other of these possible mews. I agree with Dr. Connick that both points of view should be presented for the sake of scientific completeness.
