June 20,1952 [CONTRIBUTION
SPECTROPHOTOMETRY OF
NO. 185 FROM
THE INSTITUTE FOR
RUTHENIUM(III)-DITHIO~XAMIDE SYSTEM
ATOMICRESEARCH AND
COLLEGE 1
THE
DEPARTMENT OF
CHEMISTRY,
Spectrophotometric Investigations of Some Complexes of Ruthenium. Ruthenium-Dithiooxamide System BY
3163 IOWASTATE
111.
The
RUTHPOWERS YAFFE~ A N D ADOLFF. VOlGT RECEIVED E ~ B R U A R7,Y1952
The ruthenium(1V)-dithioilxamide system has been studied spectrophotometrically. It was found that both Ru(111) and Ru(1V) form the same blue-green complexes, Ru [SC(NH)CSNHt)+*and Ru [SC(NH)CSNH*]*. The formation constants of these complexes were evaluated a t unit ionic strength. In the determination of the formulas and stability conThe reaction of ruthenium(II1) and ruthenium stants, in all solutions, the ruthenium concentration was (IV) with dithiooxamide in acetic acid to form a maintained constant at 2.793 X 10-6 M, the perchloric blue-green color has been developed into a colori- acid constant at 0.127 M, the acetic acid constant at 50%. metric method for the analysis of r u t h e n i ~ m . ~ -and ~ the ionic strength constant at 1.0. Solutions were preAlthough the reaction has been used in a colorimet- pared by mixing necessary quantities of a solution of ruperchlorate and a freshly prepared solution of ric procedure, none of the investigators has re- thenium(1V) dithioijxamide to give the final desired concentrations. ported any study of the nature of the reaction to Equilibrium was attained by the previously described heatdetermine the formulas and stability constants of ing and cooling procedure. More than sixty solutions which contained ratios of dithidxamide to ruthenium from 360 the complexes involved. Preliminary investigation of the ruthenium-di- to 1 were so prepared and scanned. thiooxamide reaction suggested that, as in the thioDiscussion cyanate? and thioureas reactions, the Ru(1V) was being reduced to Ru(II1) a t the expense of the comA typical constant wave length plot of the optical plexing agent. Therefore the method of interpre- density as a function of the dithiooxamide to ruthetation of spectrophotometric data was that pre- nium ratio is given in Fig. 1. This shows clearly sented in the previous paper^.^^^ that the optical density of the solution does not change for ratios of dithiooxamide to ruthenium Experimental greater than 140 (the line labeled ds in Fig. 1). Absorption Measurements.-A Cary automatic recording spectrophotometer was used for all the optical density measurements. Matched silica cells of 5.0O-cm. optical length were used. The results are expressed in terms of o tical loglo f i ~ / I ) density, D,defined by the relationship, D where IOand I are the incident and transmitted l g h t mtensities. respectively. Materials.-The preparation and standardization of the Ru(ClO4). solution have been given in a previous paper.' DithiMxamide, Eastman Kodak Co. Organic Chemical 4394, was weighed out directly to provide solutions of the desired concentrations. Since it is only slightly soluble in cold water, the solvent used wm a mixture of water, perchloric acid and acetic acid. The LiClO4 and He104 were from G. Frederick Smith Chemical Co., while the acetic acid was from Mallinckrodt Chemical Works. These were purified, recrystallized and standardized as necessary by standard procedures. Particular care was taken to avoid iron contamination. Technique.-As reported previously,@it was found that heating was nccwary to attain equilibrium within a reasonable length of time. A minimum period of heating of 30 min. in a boiling water-bath was required to produce equilibrium. No decomposition of the complex was noted even when heating periods as long as two hours were investigated. In this study all solutions were heated exactly 45.0 min. in a boiling water-bath, and then cooledoin ice until the temperature of the solution reached 25 f 1 Immediately after preparation, each solution was scanned as quickly as possible on the spectrophotometer. It was determined on the Cary that the colored solution, after preparation, was stable for at least two days at room temperature.
-
.
(1) Work performed in the Ames Laboratory of the Atomic Energy Commission. (2) Taken from part of the Ph.D. thesis of Ruth Powers Yaffe, Iowa State College, 1951. (3) H. WBbling, Ber., 67B,773 (1934). (4) H. Wfilbling and B. Steiger, Mikrochemk, 16, 295 (1934). (5) F. J. Welcher, "Organic Analytical Reagents," Volume 4, D. Van Nostrand Co., Inc., New York, N. Y.,1948. (6) G. H. Ayrei and F. Young, A d . C h m . , S I , 1281 (1960). (7) R. P.Yaffe and A. P,Voict. Taxa Jou~NA~., T4.2600 (1982). (0) R. P. Yallc and A, 8, Vow, Wd., T4, a608 (1919).
Q
a
Maximum Exprrimrntol Error
100 [Et]y[R"]. 120
0
20
RaYo
140
-
Fig. 1.-Optical densities at 700 mp of various solutions with excess dithiobxamide present: [Ru] 2.793 X lo-. M; [HClOd] = 0.127M; [HOAc] = 50%; p = 1.0; cell length = 5.00 cm.; circles, experimental points; curve6 calculated on the basis of Ru(dt)l+Pand Ru(dt),.
Hence the adsorption spectrum of the last complex is known from experimental data. The absorption spectrum of this complex a t various concentrations is given in Fig. 2, from which it may be seen that the complex obeys Beer's law. The method used to interpret the data was that presented in previous papers.?Bs Noting the similarity in the plot of optical density against ratio between the dithiooxamide system, and that of the thiocyanate system,' it seemed reasonable to assume that only one complex was formed. Calculations were made on this assumption, but under no circumstances could the data be fitted adequately, Therefore the assumption that only one complex wa8 prtsent appeared to be crrotleou$,
RUTHPOWERSYAFFE
3164
0.~1I
0.8
W o v e l e n p t h in
Millimicrons
Fig. 2.-Absorption spectra of Ru(dt)a; [HC104] = 0.127M; [HOAc] = 50%; p = 1.0; cell length = 5.00 M ; B, [Ru] = 2.234 X 10-6 cm.; A, mu] = 2.793 X M; C, [Ru] = 1.076 X 10-6 M ; D, [Ru] = 1.117 X 10-6 M; E, [Ru] = 5.585 X lO-'M.
Assuming that two complexes are present, the following equations may be written
+ nHdt
Ru+3 Ru(dt),,3--"
+ qEIdt
+
R ~ ( d t ) , , 3 - ~ dit Ru(dt),,+,"-"-q -b qII+
where EIdt is used as an abbreviation for dithiooxamide. The non-thermodynamic equilibrium constants then become K, = [Ru(dt),8-" ] [H+In [Ru +3 ~. I [Hdt 1" Kn+q
=
,
AND ADOLF
F.
VOIGT
Vol. 74
In each approximation, the calculations of both K1 and KS were carried out a t 12 different wave lengths a t 10-mp intervals from 640 to 750 mp and the constants averaged to give the final values. It should be pointed out that while the precision of the calculations in the dithiooxamide system is as good as in the thiocyanate7 and the thioureas systems, the accuracy is not. One reason for this is that the complexing agent is colored, making it necessary to correct all optical density measurements. Also, the complexes are fairly stable, so that a smaller excess of ligand is necessary to complex the ruthenium. The approximation that the equilibrium concentration of the dithiooxamide is essentially equal to the total analytical cqncentration then becomes somewhat inaccurate. Using the final values found for the equilibrium constants K1 and Ks, the absorption spectra of the first complex, Ru(dt)lf2, (Fig. 3) and of the uncomplexed ruthenium were calculated. Even though the solutions contained 50% acetic acid, the absorption spectrum of the uncomplexed ruthenium was essentially that of ruthenium(II1) perchlorate and differed from ruthenium(1V) perchlorate sufficiently to justify the conclusion that the ruthenium is in the I11 state in the complex.
8-n--0 ] [H 14 [ R ~ ( d t ) , , ~ *[ ]H d t ] ~
[Ru( dt ).+
+
By use of the method for two complexes already given,*the following equation may be obtained
i-I I I I I where do, dl and d3 are the optical densities of the 400 500 600 700 800 W o v r l a n g t h i n Milllrnlcrons. uncomplexed ruthenium, the first complex, Ru(dt)n3-n, and the second complex R ~ ( d t ) ~ + ~ ~ - ~ - -Fig. 9 , 3.-Calculated absorption spectrfim of Ru(dt)l +*: respectively. This equation was used in a series of [Ru] = 2.793 X 10-6 M; [HC104] 0.127 M ; [HOAC] = successive approximations in the analysis of the 50%; p = 1.0; cell length = 5.00 cm. system . Using the calculated equilibrium constants Kl It was found that the data would fit only if n is taken as one, and p as two. Thus the equations and K3 and the calculated values of do and dl, together with the experimental value of &, the optical should be written density curve as a function of dithiooxamide to R u + ~ Hdt 2 Ru(dt)l+' + H + ruthenium ratio was calculated a t 700 mp. This is Ru(dt)r 2H+ Ru(dt)i+' 2Hdt the curve drawn through the experimental points in The non-thermodynamic equilibrium constants are Fig. 1. To test the acetic acid dependence of the system the procedure was repeated a t an acetic acid No evidence was found for a two-to-one complex, concentration of 75% instead of 500/00. It was found that the reactions themselves are independRu(dt)p+. Four successive approximations were carried out ent of the concentration of acetic acid present, alfor the evaluation of K , and Ks. The results of though changing the acetic acid concentration alters the solvation of the ruthenium. these calculations are summarized in Table I. It should be pointed out that all of the equations TABLE I and equilibrium constants in the ruthenium-dithioSUMMARY OF SUCCESSIVE APPROXIMATIONS TO K1 AND K s oxamide system have been written as a function of Approximat ion Ki K1 the hydrogen ion concentration. To test this de1 1400 f 400 12000 f 600 pendence, the procedure was repeated a t 0,254 M 2 990 f 90 9300 f 300 perchloric acid, instead of 0.127 M. This work was 3 930 f 30 8400 f 300 not done in great detail since the nature of the re4 aao rc 30 8300 f 200 action was now known, Only the first approxima-
+
+
+
SOLUBILITY OF WATER IN LIQUID PHOSPHORUS
June 20, 1952
tion was carried out in the calculations. This was compared to the first approximation calculation at the lower acidity. The results of these measurements and calculations are summarized in Table 11.
3165
ruthenium is held tightly in the complex. This may be illustrated by obtaining the formation constants for the reactions Ru+3
+ dt-
I_ Ru(dt)i+2
TABLE I1 COMPARISON OF FIRSTAPPROXIMATIONS TO K1 AND Ks IN 0.127 M AND 0.254 M HClOi (a) Assuming hydrogen ion dependence
0.127 0.254
12000 f. 600 11000 f 3 0 0 0
1400 f 400 1300 f 60
(b) Not assuming hydrogen ion dependence [Ru(Hdt)i +'I [Ru(Hdt)a+a] [HC1o'l' M
K 1 = [Ru+a][Hdt]
0.127 0.254
11000 i 3000 5200 i 200
Ka = [Ru(Hdt)i+a][Hdt]~
740000 f 32000 180000 i40000
From Table IT, it is seen that the assumption of hydrogen ion dependence gives agreement in both equilibrium constants at the two acidities, whereas the assumption of no dependence does not. Therefore it must be concluded that dithiooxamide is behaving as an acid, each molecule releasing a proton on formation of the ruthenium complex. This behavior was anticipated since dithiooxamide is known to be a weak monobasic acidas Although the formation constants for the reactions as they have been written are quite small, the (9)
R.P.Yaffe and A. F. Voigt, THISJOURNAL, 74,2941 (1952).
[COYTRIBTJTION FROM
THE
Using the acidity constant of dithiooxamide evaluated a t unit ionic strength: k, = (3.78 f 0.04) X lo-", K1' and K3' were calculated. K1t = K_I (9.3 f 0.3) X IO* = (2.4 f 0.1) X IO'* I
ka
(3.8 f 0.1) X 10-11
Dithiooxamide has bem reported to form fivemembered [C-C-N-S-Metal] chelate rings with both platinum and p a l l a d i ~ r n . ~ -It ~ seems quite likely that the ruthenium-dithiooxamide complexes involve chelate rings of the same type (I). H N==C--S-Ru
1
S=C-N
..*
/3
H2 '
I
AMES,IOWA
DEPARTMENT OF CHEMISTRY, ~JNIVERSITY OF CALIFORNIA, BERKELEY ]
The Solubility of Water in Liquid Phosphorus BY G. L. ROTARIU, E. W. HAYCOCK AND J. H. HILDEBRAND RECEIVEDJANUARY 21, 1952 Determinations of the solubility of water in liquid phosphorus gave 3.6 f 0.3 mg./g. a t 45" and 3.9 f:0.5 a t 25". We select 6.7 X 10-3 ml. of HzO/l ml. P4 a t 25" for evaluation. Published figures for He0 in C& and CCL give 1.95 and 1.59 ml. of HzO/rnl. of solvent, respectively. The solubility parameter of water calculated from these data gave 26.2 for HzO in P4,26.1 in CSZand 24.7 in CCla. A previously published analysis of data for HzO in paraffins gave values close to 24.5.
In several recent investigations involving white phosphorus' we have found it far more convenient to protect it from air by a layer of water than by operating in a vacuum, and we have had satisfying evidence that the properties under investigation were not appreciably affected by dissolved water. We decided, however, that it would be desirable to know the solubility of water in phosphorus, both for its bearing upon past and future experiments, and for its theoretical interest, in view of the extraordinarily high solubility parameter of phosphorus and the dipole moment of water. The procedure adopted was, briefly, to equilibrate liquid phosphorus with water, freeze it rapidly, presumably entrapping dissolved water in the solid as minute droplets, as in the case of dissolved mercury; wash off all exterior water by repeated application of cold dry methyl alcohol; melt the phosphorus and shake ufider methyl alcohol to
The essential details of the procedure were as follows. The phosphorus used in determinations 1-7 and.lG-18 was purified by treatment with dilute acid dichromate, as previously described; that used in determinations 8-15 was distilled i n vacuo. The methyl alcohol was Eimer and Amend C.P., anhydrous, handled so as to prevent access of water from the air. The Fischer reagent was prepared and standardized carefully by Mr. Tashinian, of our microlaboratory, to whom we express our grateful appreciation. Five ml. of purified phosphorus was pipetted under water into a previously tared flask with ground glass cap. This was shaken in a water-bath at a stated temperature. One hour was found amply sufficient for equilibrium. The flask was then quickly chilled t o freeze the phosphorus, care being taken to keep it in one compact mass and not entrap drops of the surrounding water. It was detached from the glass, and the flask cooled in ice-water, to reduce the amount of dissolved phosphorus, which can affect the reagent. The water in the flask was replaced by six washings with 50-ml. portions of precooled methyl alcohol. A seventh 50 ml. was allowed to remain in the flask. The sixth portion was
(1) R. E. Powell, T. S. Gilman, G. J, Rotariu. Eva Schramke and J. H. Hildebrrnd, TEISJOURNAL, 71, 2626, 2627 (1861).
(2) Cf.J. Mitchell, jr., and D. M. Smith, "Aquamctty," Interscience Publimheta, Inc,, Nos Voork, N. Y , , 1848, p, 18,
extract the water, and determine its amount by the Fischer titration method.2
