16
HANSl3. JONASSEN,G. G. HURW,B. B. LEBLANC AND A. W. MEIBOHIM
VOl. 56
INORGANIC COMPLEX COMPOUNDS CONTAINING POLYDENTATE GROUPS. VI. FORMATION CONSTANTS OF COMPLEX IONS OF DIETHYLENETRIAMINE AND TRIETHYLENETETRAMINE WITH DIVALENT IONS' BYHANS B. JONASSEN, G. G. H U R S TR. , ~ B. LE BLANC,^^ AND A. W. MEIBOHM~~ Richardsan Chemical Laboralories, Tulane University, New Orleans, Lozci9aknu Raesived August SO, 1961
The formation for the ions Mn(II), Fe(II), Co(I1) and Zn(I1) have been measured with diethylenetriamine and triethylenetetramine a t 30" and 40". Formation constants and other thermodynamic properties have been calculated for them complexes. The order of stability of these complexes is related to the second ionization potential of the metal.
The authorsa*'previously measured the formation curves and determined the formation constants of the complexesof diethylenetriamine (abbr. dien) and triethylenetetramine (abbr. trien) with nickel(I1) and the copper(I1) ions. This paper reports the determination of the formation curves and formation constants of the complexes of these two polyamines with the manganese(I1)' the iron(I1) the cobalt(I1) and the rinc(I1) ions at 30 and 40' and some of the thermodynamic values calculated from these data. The method used for the measurements is the pH method develoDed bv J. B j e r r ~ m . ~The specific equations for the complexes-of dien and trien were developed previously.%4 From the stepwise acid-base dissociation constants of dien H,+3and trien H F and the pH's of solutions .containing known concentrations of the metal ion, a mineral acid, and dien or trien the concentrations of the free amine and of the complexed amine can be calculated. From these the values of n, the mean number of amine molecules coordinated to each metal ion, and p(dien) or p(trien), the negative logarithm of the free dien or trien concentration, can be computed. These quantities substituted into the following equations, yield %he values of the formation constants of the complex ion Me(dien) +2 and Me(dien)Z2. -
-1nu-
n
-
-1 + log : T + ._
log k, = p(dien)
log
n L1 4- (n - I)[dion]kl]
Similar equations5 huve been derived for the com-
plex ions Me(trien) +* and Mez(trien)z4
(1) Presented at the Syinposiuln on Complex Ions and Polyelectrolytes. Ithaca, New l'ork. June 18-21, 1951. Abstracted from the Ph.D. dissertations of R. B. LeBlanc and A. W ' . Meibohm, February, 1950.
(2) Present addresses: (a) Department of Chemistry, Mississippi lout hern College, Hattiesburg, hliss. (b) Department of Chemistry,
Texas Agricultnrnl and Merhanicnl College, College Station, Texas. (e) Department of Chemistry, Valparaiso University, Valparaiao. Indiana. (3) H. B. Jonassen. R. B. LeBlanc and R. M. Rogan, J . A m . Chem. Soe.. 79, 4968 (1950).
(4) H. B. Jonasaen and A. W. Meibohm, T A I JOURNAL, ~ 61, 726 (1931).
(5) J. Ujerruiim. "Metal Ammine Forination in Aqueous Solution," 1'. Hasse a i d Son, Copcnlmagcn. 1041.
-
+ log -1 -nn n-1 log kr:t = p(trien) + log 1.5 - ii log kl = p(trien)
The log k values c m also be obtained from the formation curves a t E = 0.5 and 6 = 1.5 for the dien complexes and 5 = 0.5 and 6 = 1.25 for the trien complexes. The partial aminium complexes which Schwarzenbach obtained by modification of B j e m ' s method could not be identified under the conditions of this investigation.' Spectrophotometric investigations of the metalamine solutions under the various conditions given in Schwarzenbach's work showed that only the density values of the absorption charachristics change with change in conditions. It seems unlikely that the amine and aminium complexes have the same absorption characteristics. Experimental Stock solutions of all the metal ions except the iron(I1) ion were prepared by dissolving a mole of the C.P. nitrate salt in a liter of solution. These solutions were standardized b electrodeposition and/or precipitation as the pyrophosp{ate: Boiled distilled water was used for these solution^. A rapid stream of nitrogen was passed through the cobalt(I1) and the manganese(11)solutions for ten hours to ensure the absence of appreciable quantities of dissolved oxygen. A stock solution of iron(I1) ion w8s prepared by dissolving 5.6 g. of C.P. iron powder in about 60 of 6 N HCl. The solution was filtered and diluted to a hter with boiled distilled water. Dissolved oxygen and any trace of iron(I1) were removed by bubbling hydrogen through the solution for twelve hours in the presence of few drops of colloidal palladium and indigo tetrasulfonate as catalysts. The iron was determined gravimetricallyas the oxide, the chloride as AgC1, and the excess HCl by potentiometric titration in an atmosphere of nitrogen. A 2 M KNOJ stock solution and a 4 M KC! stock solution were prepared by weighing the dry salts and hasolving in the required amount of solution. A 1 M "01 solution was repared from boiled C.P. HN03 and standardized with heagent Grade N&COo. Technical grade dien or trien from Carbide and Carbon Chemicals Corporation was purified as previously described-7 It was dissolved in a liter of solution, and titrated potentiometrically with the standard 1 molar mjd. The solutions (except for that of the mn(I1) ion) used for the measurements were prepared in 100-ml. volumetric flasks. They contained 0.1 M of metal ion, 0.1 M of "03, 1.OO M of KN03 and KCl, and varying concentrationsof the amines. The 1 molar stock solution of the amines also was 1 molar in KCl. The solutions for the iron(I1) ion measurements were
e.
(6) G. Schwarzenbach, H d u . Chim. Aclu, 38, 974 (1950). (7) H. B. Jonassen. R. B. LeRlanc, A. W. Meibohm and R. M. Rogan, J . Am. Chem. Soc.. 74,2430 (1980).
Jan., 1952
FORMATION CONST~LNTS OF DI- AND TRIETHYLENETETR~LM~INE WITH DIVALENTYION~ 17
0.01 M in iron(I1) ion, 0.01M in HCl, 1.27M in KCI, and of varying concentrations in amine. In order to minimhe 0x1dation HCl and KC1 were substituted for HNG and KNOI, and the concentrations of the imn(II) ion and the acid were decreased to 0.01 M. The concentration of the KCl was increased to 1-27M to make the salt mnmtration the same as that used for the measurements for the acid-base constanta of the amines and the formation constsnts of the other
metal complexes. The solutions of the cobalt(II) ion, the iron(II) ion and the manganese(I1) ion were prepared up to the point of addition of the stock solution of the metal ion. Nitrogen was bubbled through the solutions in the volumetric flasks for 45 minutea to remove dissolved oxygen, whereupon the metal ion solution was added. This recaution was taken since these three ions are easily oxi& in bsliic solution by digsolved
TABLE I1 FORMATION CONSTANTS OF MANQANESE(II) AND DIETHYLENE-NE AT 30"
Results
4.80
1
4.29 4.19 3.97 3.80 3.66
.+
Wien
0.0450 .0551 .Of355 .OS50 .0900 .lo47 .1179 .1302 .1451 .16W .1901 .1999 .2198 .2402 .2598 .2850
5.25 7.52 7.84 8.10 9.15 8.26 8.35 8.42 8.52 8.62 8.73 8.88 0.01 0.20 9.40 9.62
9.15 4.44 3.80 3.29 3.19 2.97 2.80 2.66 2.47 2.28 2.07 1.80 1.55 1.24 0.91 0.58
2.25 0.0999 1.98 .loo1 .loo1 1.96 1.93 .loo1 1.92 lo00 1.90 lo00 1.88 .0999 .loo2 1.86 1.83 .0999 1.79 .0995 1.73 .0996 1.65 .ms 1.56 OB98 1.40 .loo2 1.20 .W96 0.96 .0996
Wien
m c
.
.
.
-
.05 .09
- .16 - .28
-
-
A = log (1
n + (5- l)[dien]ki
-n 1.099 1.293
3.28 3.07
-0.96 - .38
1.468 1.640
2.80 2.55 2.24
+ .25 + .54
1.777
- .06
2.82 2.88 2.82 2.84 2.80
+0.50
+ .19 + .08 + .04 + .02
-
Table I11 shows the values of the formation constants at 30" and 40°, the free energies of binding, and the heats of binding of the complexes of dien with the four ions. The data for nickel(I1) ion and the copper(I1) ion which were previously reporteda are included for comparison.
TABLE I11 VALUES FOR COMPLEX IONS OF DIETHYLENETRIAMINE AND DIVALENT METALI O N S All values in kcal./mole; A H values = AHaao THERMODYNAMIC
-log
dien
4.05 3.98 3.96 3.98 3.97 3.99
Average 2.83
pH MEASUREMENTSOF MANQANESE(II) SOLUTIONS CONTAININQ D~ETWLENETBUL~~~NE AT 30" CMN(II, 0.0949M. CKNO~ CKCI= 1.oOK; C. = total concentrahon of hydrogen ions bound to dien; = total concentration of dien in solution; C'dian total concentration of dim not bound in a complex ion; %inn = averaw number of hydrogen ions attached to one uncomplexed dien molecule. PH
log ki
-
+ + +
TABLE I
Cdian
A
-0.01 - .04
n
-0.74 - .27 - .18 .09 .33 .61
Average 3.99
balt(I1) ions during the measurements.
Table I contains the data and calculations on the manganese( 11) ion and dien a t 30'. The two formation constantsof the manganese(I1) ion and dien complexes a t 30" calculated from points on the formation curve are shown in Table 11.
log
p(dien)
0.153 .349 .400 .549 .683 .804
o T t i H ' s of all the solutions were measured at 30 f 0.2" and 40 f 0.2'. A Beckman model G H meter, standardized with Beckman buffer solutions (& of I 4, % 7 and 1.0) was used. An atmasphere of nitrogen was maintained over the solutions of the manganese(II), the iron(I1) and the co-
-+ -
n
C'dien
n
0.0443
0.00s
.0507
.046
-0510 .0519 .0520 .0526 .0531 .om9 .0547 .0557 .0574 .0606
.153 .349 .4oo .549
.0642
.0718 .os35 .lo35
.683 .804
.953 1.099 1.293 1.468 1 A40 1.777 1.859 1.913
P(dien) 10.15 5.44 4.80 4.29 4.19 3.97 3.80 3.66 3.47 3.28 3.07 2.80 2.55 2.24 I .91 1.58
Figurc 1 shows thc formation curves at 30" for dien and the iron(TI), the cobalt(II), the einc(I1). manganese(II), coppcr(I1) and nickel(I1) ions, respcctivety.
Fig. 1.-Formation curves of diethylenetriamine and the followingions a t 30": Mn,U;C!a -0-; Cu,&; Fe, -A-; Ni, -A-; Zn, -0-.
Cation Mn(I1) Mn(I1) Fe(I1) Fe(E1) Co(I1)
Co(I1) Ni(I1) Ni(l1) Cu(I1) Cu(I1) Zn(I1)
Zn(II)
T. O C . log kl log kr 30 40 30 40 30 40 30 40 30 40 30 40
3.99 3.89 6.23 6.03 8.47 8.26 10.81 10.54 16.11 15.64 9.14 8.95
2.83
2.72 4.13 3.95 6.07 5.83 8.14 7.83
--PI 5.5 5.6 8.6 8.6 11.8 11.8 13.0 15.1 22.3 22.4 12.7 12.8
-AP: 3.9 3.9 5.7 5.7 8.4 8.4 11.2 11.2
-AH,
-A& 5 8
lo 12
13
20 8
Figure 2 shows thc formation curvcs of the complexes of triethylenctetramine and divalent mctal ions.
I
"'t
*triad.
Fig. 3.-Formation
curves: complex ions of all cations a t 30".
18
HANSB. JONASSEN,G. G. HURST, R. B. LEBLANCAND A. W. MEIBOHM
Vol. 56
Table I V gives the thermodynamic values obtained for the complex ions of triethylenetetramine with the above divalent metal i:ns as well as cadmium(11) ions at temperatures of 30 and 40 TABLE IV THERMODYNAMIC VALUESOF COMPLEX IONSOF TRIETHYLENETETRAMINE AND DIVALENT METALIONS All values in kcal./mole, AH values =
the case in the copper and zinc ions and only the first formation constant for these complexes, therefore, can be obtained from these data. . Figure 2 shows that in each case the formation of a complex ion of triethylenetetramine with one metal ion occurs. In the presence of excess amine expansion of the coordination sphere occurs for all but the copper(I1) and zinc(I1) ions. Cation T, 'C. log kl log I:* -AF, - A h : * -AH1 -Ha:* Comparison of the data of cadmium(I1) and Mn(I1) 30 5.43 2.84 7.5 3.9 4 zinc(I1) shows that the 1 : l complex ion is more Mn(1I) 40 5.31 2.72 7.6 3.9 Fe(I1) 30 8.31 3.92 11.5 5.4 stable for zinc(J.1) than for cadmium(I1) and that 9 Fe(I1) 40 8 . 0 8 3 . 7 0 11.6 5.3 the coordination sphere of zinc(I1) is not expanded COW) 30 11.21 3 . 3 8 1 5 . 5 4.7 9 4 to a coordination of six whereas that of the cadCo(I1) 40 11.03 3 . 2 7 15.8 4.7 mium(I1) ion is. It is to be expected that the . Ni(l1) 30 14.34 5.63 19.8 7.8 13 9 Ni(I1) 40 14.01 5 . 4 1 20.1 7.8 larger cadmium(I1) would act in this manner. Cu(I1) 30 20.62 28.6 The equations of the method of this paper require 22 Cii(I1) 40 20.08 28.8 that all nitrogen atoms of triethylenetetramine be Zn(I1) 30 11.94 lR..i 4 coordinated to the central metal ion. The coordiZn(I1) 40 11.81 10.9 Cd(I1) 30 10.92 :Ll9 13.1 4.4 nation number of six thus requires a formula [ M Z 4 Cd(I1) 40 10.79 3.07 15.3 4.4 A,-. 1+4. It is possible to postulate other formulas for these Discussion complexes such as MA2 where two of the nitrogen Formation of Complexes.-As shown in Table atoms of one of the amines would not be coordinated. I11 the manganese(II), the iron(II), the cobalt(I1) However, spectrophotometric studies8 and polaroand the nickel(1I) ions exhibit a coordination graphic studies of several of these complexes indinumber of six coniplexing with two dien molecules. cate that the complex ion with the highest amine The copper(I1) ion and the zinc(I1) ion exhibit a to metal ion ratio is the 3/2 : 1 complex, coordination number less than six, but still coordiThe steep increase in the formation curves of the nate two dien molecules. As was explained previ- cadmium(I1) and manganese(I1) complex ions with ously3 the equations for the calculations of forma- trien shown in Fig. 2 is no doubt due to the fact that tion constants are applicable to these ions for an ii a t these high concentrations of amine polymeric value greater than one only if all the amino nitro- hydroxy complexes may exist which will change the gens are coordinated to the metal ion. This is not pH and thereby the 6 and p(amine) value. Such polymers would not allow coordination of all nitroI gen atoms in the amines and Bjerrum's method would not be valid. A previous polarographic investigation of the complex ion formed between cadmium(I1) and triethylenetetramine9 gave log k values of 13.9 for the 1 : l complex ion. However, the log k values for this complex compound obtained in t,his investigalion is 10.9, arid the value for the 3 : 2 complex ion was estimated to be 3.1. The difference in these values of the 1:1 complex ion can be expected since a large excess of the amjne (100: 1) and greater is present in polarographic investigation whereas the excess of amine in the solution from which these data were obtained was never larger than 2 : l . At such small excess the above authors found that the electrode process was irreversible indicating the possible presence of other comples ions. The complex (Cdz trien,) f 4 found in this inilestigation is in accord with srich coiiclusions. Order of Stability of Ions.-The order of stability of the complex ions of the same metal ion is that expected from the nature of the amine. The stability of the ions formed by PP'P"-triaminetrjethylamine (tren), an isomer of triethylenetetramine indicated that steric effects are important.6 It is st'ericnlly impossihle for this amine to enter a planar configuration as required by the copper(I1) ion and the decrease in stability, relative to the triFig. 3-Rclntivc stnbility of coml)lcs ions, coiirdiriatiori (8) 1.1. D. Joriwscri nitd n. E. Dottglns, J . .4?nn.Chem S n c . , 71, 4094 nuiul)er of four: et,liyleiictdiamiir~,--A-; propylene- (1049). diet,hglcnetrinmine,-0 -; triet hylcnediamine, -@-; (9) 13. 1,;. l)niigIas, tetramine, --; ~ , ~ ' , ~ " - t ~ r i n ~ r i i ~ i n t ~ -@-. i ~ t l ~ ~ l ~2484 ~ i ~(1920). ii~e,
11. .\. 1.niliiteit sild J. C'. Bailar, Jr.. i b i d . , '72,
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 complex FeIf2; Mn(C20a)+ yields manganous transfer have come some promising possible gener- very much more rapidly than does Mn(CzO&-. alizations. One of these might be stated as folThe present work examines the Cu(I1)-CNlows: coordination complexes are involved mecha- reaction from the points of view outlined above. nistically as intermediates in homogeneous ionic oxidation-reduction reactions. This hypothesis Experimental appears to be particularly applicable to reactions Reagent grade chemicals were used. Acidified cupric involving cationic oxidants and anionic or “Lewis- chloride stock solution was standardized through thiosulfate base” reductants. For example, systems in which against KpCraOTwith a small quantity of Na2COaadded, and such complexes have been observed as intermedi- potassium cyanide was standardized against AgNOI with the The ammonia was cooled, diluted, ates are Fe(111)-I -, Mn (111)-C2Or, 2, Ce (1V)- iodide end-point. against HCI, and thereafter kept in an ice-box, glyc01,~ Ce(IV)-CI-,5 Fe(III)S03-,6 Ce(1V)- standardized as were all solutions containing ammonia. Kinetic runs CzOa,7 and a number of others. were made at 0’ f 0.lo, maintained by a bath of melting A point of further interest in connection with ice in a dewar flask. Scparate erlenmeyer flasks containing 50 ml. of the dethese coordination intermediates is any relationship cupric ammonia solution and a cyanide solution of apwhich might exist between the number of oxidizable sired propriate concentration were cooled to equilibrium in the iceanions in the complex and the rate of “electron bath. Five ml. of the cyanide solution were then added to transfer.” For instance, the complex Fe12+yields the rapidly shaken copper solution by means of a pipet calito give 5.00 f 0.02 ml. Five-ml. portions of the referrous ion very much more rapidly than does the brated acting solution were removed by a calibrated pipet at known (1) A. V. Hershey and W. C . Bray, J . A m . Chem. Soc., 08, 1760 times and quenched in 2 ml. of a solution containing 0.2 M (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).
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-
