694
VINCENTL. HUGHES AND ARTHUR E. MARTELL
VOl. 57
. SPECTROPHOTOMETRIC DETERMINATION OF THE STABILITIES OF
ETHYLENEDIAMINETETRAACETATE CHELATES BY VINCENTL. HUGHES AND ARTHUR E. MARTELL Contribution of the Department of Chemistry of Clark University, Worcester, Mass. Received January 29, 1065
The interactions of colored metal chelates of ethylenediaminetetraacetic acid with various metal ions were determined by a direct spectrophotometric method in nitrate and perchlorate media. Under conditions such that the metal chelates are not appreciably dissociated and the chelate containing one metal ion per ligand was the important form, the measurements allowed calculation of the relative stabilit constants of the metal chelates investigated. The dissociation of the nickel(I1)eth lenediaminetetraacetate chelate was Lollowed spectrophotometrically, the absolute value of the stability constant of this cherate was calculated, and the relative values for the other metals were converted to absolute stability constants. Studies were also made of the effect of ionic strength on the stabilities of the chelate compounds.
Although it is not possible to evaluate directly from potentiometric measurements of pH, the very large chelate stability constants for the transition and heavy metal ethylenediaminetetraacetates, a number of indirect methods have been developed in recent years. Schwaraenbach and Freitag‘ determined potentiometrically the position of the equilibrium in a system containing magnesium and transition metal chelates of P,P’,P”-triaminotriethylamine and ethylenediaminetetraacetic acid. Also, the equilibrium between the ferrous and ferric chelates, determined by an oxidation-reduction method by Schwarzenbach and Heller12 and by a polarographic method by Kolthoff and Auerbach3 .has been useful in the calculation of chelate stabiIity constants. Perhaps the most unusual method is that of Jones and Long,4 wherein the kinetics of exchapge of ferric ion with the ferric ethylenediaminetetraacetate chelate was studied. Extrapolation to zero time made possible the calculation of the concentration of free ferric ions, and, from this quantity, the chelate stability constant. In recent publications,5~6 a direct spectrophotometric method for the determination of relative stabilities of transition and rare earth metal chelates of ethylenediaminetetraacetic acid has been described in principle. However, the buffering agents employed to control pH also formed weak complexes with the metals being investigated. This method has now been investigated for solutions with no interfering buffers, and with an ionic atmosphere controlled with nitrates or perchlorates. Further, the relative chelate stability constants thus obtained have been placed on an absolute basis by spectrophotometric measurement of the degree of dissociation of the nickel chelate in sufficiently acidic solutions. The effect of ionic strength on the relative stability Constants also was investigated. Experimental Part Materials and Equipment.-Absorption measurements were made with a Beckman model DU quartz spectrophotometer with 10.003-cm. cells. A Beckman Model G pH meter with extension electrodes was used to measure p H . (1) G. Schwarzenbach and E. Freitag, Helu. Chim. Acta, 34, 1503 (1951). (2) G.Schwarzenbach and J. Heller, ibid., 34, 576 (1951). (3) I. M. Kolthoff and C. Auerbach, J . Am. Cham. Soc., 74, 1452 (1952). 56, 25 (1952). (4) S. S. Jones and F. A. Long, THISJOURNAL, (5) a. C. Plumb, A, E. Martell and F. C. Bersworth, ibid.. 54, 1208 (1950). (6) A. B. MWtell rand M L 0. f l u m h , ibrd., 66, 993 (1962).
It was calibrated with standard acetate buffer according t80 the method of Michaelis.? The ethylenediaminetetraacetic acid employed was furnished through the courtesy of the Bersworth Chemical Company, Framingham, Mass., and was further purified by two successive recrystallizations from water. The metal nitrates were Baker’s Analyzed and Mallinckrodt C.P. products. The metal perchlorates were prepared from the corresponding nitrates by precipitation as the metal hydroxide which was then washed by decantation with distilled water until free of soluble salts. The product was then dissolved in perchloric acid and the p H was adjusted to between 4.0 and 4.5 with perchloric acid. The metal salts were standardized by a metal chelate titration developed by Martell and Jones.8 The potassium nitrate, potassium perchlorate and perchloric acid were J. T. Baker Analyzed grade reagents. The label analysis listed maximum impurities to be less than 0.02% in each case. Carbonate-free potassium hydroxide was prepared by a modification of the method of Schwarzenbach and BiedermannP A 0.1 M potassium chloride solution was allowed to react with two successive batches of silver oxide over an extended period of time with occasional shaking. The resulting potassium hydroxide solution contained no detectable amount of chloride ion. All solutions were brought to equilibrium in a constant temperature bath at 30” before absorption measurements were made. The temperature of the absorption cell in the spectrophotometer was found to remain a t 30 i= 1”. The metal complex was prepared by mixing solutions of the metal nitrate and the potassium salt of the complexing agent, and the pH of the resulting solution was adjusted to 4.5. For the determination of relat,ive stability constants, solutions were made up which contained equimolar amounts of each of two metal ions and of the dipotassium salt of ethylenediaminetetraacetic acid. The pH of the resulting solution was considerably lower than the solutions used because of displacement of hydrogen ions from the complexing anion. These conditions were employed because it was found that equilibrium was obtained relatively rapidly a t low p H , whereas a similar solution prepared in the pH range 4-5 required in certain cases a number of days before equilibrium was reached. After the acid solution had been allowed to remain in the constant temperature bath for a few hours, the pH and ionic strength were adjusted to the desired values. The solutions were checked for equilibrium’by determination of absorption over a period of time a t various concentrations. All mixtures were found to have reached equilibrium after standing in the constant temperature bath overnight. The Cu(I1)-Pb(II), Ni(I1)-Pb(I1). and Cu(I1)-Ni(I1) systems were studied as a function of ionic strength in potassium perchlorate while the first two were also studied in potassium nitrate. Calculation of Relative Stabilities.-The solutions of the metal nitrates and perchlorates, the metal complexes, and mixtures of two metal salts with the complex of one of them, (7) L. Michaelis, “Physical Methods of Organic Chemistry,” Vol. I, second edition, Interscience Publishers, Inc., New York, N. Y., 1950, p. 1727. (8) L. A. Jones, Thesis, Clark University (1952). (9) G, Schwaraenbach and W. Biodermann, Helv. Chirn. Acto, 31,338 (1948)~
c
DETERMINATION OF ETHYLENBDIAMINETETRAACETATE CHELATION
Oct., 1953
at equilibrium, were all found to obey Beer’s law. Hence, the optical density of a mixture in terms of its components may be expressed by D = ecl e’c’l eacaE ea‘ca‘l (1) where c and c’ are the concentrations of uncomplexed metal ions M and M’, while ca.and ca’ are the concentrations of the corresponding metal complexes, e represents the corresponding extinction coefficient, and 2 is the length of the light path. Since it has been shown126J0 that, under the reaction conditions employed, only the 1:1 metal chelates of ethylenediaminetetraacetic acid need be considered, the equilibrium constant may be expressed by
+
+
+
where [] represents molar concentration, HIY represents ethylenediaminetetraacetic acid and M represents the metal ion. Ionic charges are omitted for convenience. It is apparent that K , = K/K‘, the ratio of the equilibrium formation constants of the respective metal chelates. The total concentrations of metal species, CM and CM’, and of the complexing agent, CY, are also known, hence
Equations ( I ) , (3), (4) and (5) are sufficient to calculate the only four unknowns, c, c‘, ca and ca’, from the experimental measurement of optical density at one wave length. These quantities were then used to calculate K , from equation (2). The values of e, e’, ea and e,’ were obtained from absorption measurements on the pure substances. Absolute Stability Constants.-In strongly acid solutions the summation of the major forms of ethylenediaminetetraacetic acid present under the reaction conditions may be expressed by ZY = [HdY] [HaY-’] [HZY-Z] =
+
+
where kl, kl, k3 and k4 are the successive acid dissociation constants of H4Y. If this equation is solved for [Y-4] and substituted into equation (3), the expression for the formation constant becomes
The dissociation constants kl, kz, k ) and ka have been determined previously by a number of workers.11-13 The values given by Carini and Martell, IC1 = 4.3 X 10-3, kz = 1.6 X 10-8, k3 = 5.6 X lo-’ and kd = 1.8 X IO-lO, were used in the calculations. These values introduced some error since they were obtained at 25” and in 0.1 M KCl solution. The values listed above required correction for hydrogen ion activity, assuming that the value reported by Harned and Hamer14for HC1 in KC1, y f = 0.781 holds for hydrogen ion. Thus the “hybrid” constants were converted to the “concentration constants” needed in the present investigation. Equation (7) was applied to the nickel(I1) and copper(I1) chelates. Appreciable dissociation was found to occur only when the hydrogen ion concentration was greater than 0.01 M . Solutions of the metal chelates, initially a t ionic strength of 0.1, were diluted with various quantities of 0.1 M potassium perchlorate and 0.1 M perchloric acid. The optical density was then taken at four wave lengths, and the concentrations of the metal complex and free metal ion were calculated with the previously-measured values of extinction coefficients of these substances. The extinction coefficients (10) See also a similar conclusion for the Nd(II1) chelate by T. Moeller and Brantley, J . Am. Chem. Soe., 72, 5447 (1950). (11) F. Carini and A. E. Martell, kbid., 74, 5745 (1952). (12) G. Sohwarzenbach and H. Ackermann, Helv. Chime Acta, 30, 1798 (1947). (13) M. J. Cabell, A. E. R. E. Report C/R813, Ministry of Supply, Harwell, Berks., England, 1951. (14) H, Hasned and W. Homer, J . Am, Cham, Soa, 66, 2194 (19833,
695
of the metals and metal chelates were checked &s a function of p H before being used in the more acid region. Neither the metals nor the metal complexes showed any change in absorption spectra from pH 7 to the point where dissociation began.
Experimental Results Relative Stabilities.-Absorptions of the free metal ions and metal complexes were measured for copper(II), nickel(I1) and cobalt(II), respectively. In all cases strict adherence to Beer’s law was observed in the spectral region from 400 to 1000 mp. The corresponding extinction coefficients were then employed in the calculation of the concentration of the species present in the equilibrium mixtures of metal pairs and ligand. The copper(11) complex was most useful for this purpose because of its intense absorption band a t 700 mp, However, in many cases it was necessary to employ the weaker cobalt(I1) bands, since the very stable copper(I1) and nickel(I1) complexes do not interact sufficiently with certain metal ions. Data for the system copper(I1)-nickel(I1)copper(I1) complex-nickel(I1) complex, and for the system copper(I1)-lead(I1)-copper(I1) complex and lead(I1) complex are shown in graphic form in Fig. 1. It is not possible to present data obtained for mixtures in the form used for solutions containing a single species. Therefore a “formal” extinction coefficient, e,, defined as follows, was employed em = D/lZci
where D = optical density I zci
= =
length of light path total concentration of one of the colored metal ions in both the complexed and uncomplexed forms
In the first region of Fig. 1, X 700-760 mp, the copper(I1) complex shows its maximum absorption, while neither nickel(II), lead(I1) ions, or the corresponding complexes, show any appreciable absorption. The formal extinction coefficient in this case is based on the total Cu(I1) concentration. On a qualitative basis, these data indicate that the nickel(I1) chelate is more stable than that of lead-
60
40
k
70
,p
25
9
2
h
x.
Fig. 1.-Formal extinction coefficient referred to the total Cu(I1) concentration: (A) Cu(I1)-Pb(I1)-EDTA acid in the ratio 1: 1:1: (B) Cu(I1)-Ni(I1)-EDTA acid, 1 :1 :1; (C) Cu(I1)-Ni(I1)-EDTA acid, 1 :2: 1; (D)Cu(XI)-Pb/II)EDTA acid, 2;3;1,
VINCENT L. HUGHES AND ARTHURE. MARTELL
696
VOI. 57
(11), since it displaces more copper(I1) from the copper(I1) complex, as evidenced by a greater lowering of the formal extinction coefficient of the copper species present. In the second region of Fig. 1, where the niclrel(I1) complex absorbs as strongly as the copper(I1) complex, the relative stabilities are not apparent from absorption data. In this case various ratios of CuY/NiY give approximately the same optical density.' In Fig. 2, data are given for mixed complex systems of nickel(I1)-zinc(II), nickel(I1)-lead(II), and nickel(I1)-cadmium(II), the formal extinction coefficients being calculated on the basis of total nickel content. The relative stabilities of the metal complexes for these metals is readily seen to be lead(I1) >> zinc(I1) > cadmium(I1).
12
64
.
Fi 3.-KO for the ratio of Cu(I1)-EDTA to Pb(I1)ED#A ions in potassium nitrate and potassium perchlorate at 30".
0
380
420
640
960
920
x.
1000
1040
Fig. 2.-Formal extinction coefficient referred to total Ni(I1) ion concentration: (A) Ni(I1)-Cd(I1)-EDTA in ratio of 1:2:1; (B) Ni(I1)-Zn(I1)-EDTA, of 1:2:1; (C) Ni(I1)-Pb(I1)-EDTA, of 1:1 :1; (D) Ni(I1)-Pb(I1)EDTA, of 2:3:2.
Not shown are the formal extinction coefficients based on the total Co(I1) concentration after equilibration of the cobalt(I1) complex with various proportions of lead(II), cadmium(I1) and zinc(11) ions. Although the relationships were not as clear in this case, a visual comparison of the curves at 460 mp, the cobalt complex maximum, indicated an increase in the interactions in the order cadmium(11) < zinc(I1) < lead(I1). I I I I I I I' The results of the calculations of K , by the 0.2 0.4 0.6 method outlined above are listed in Table I. The ai. relative stabilities a t 30' in 0.1 M KNOI of the 1:1 Fig. 4.-Upper graph, KOfor the ratio Ni(I1)-EDTA to ethylenediaminetetraacetate chelates are : Cu(I1) Pb(I1)-EDTA ions in potassium nitrate and potassium > Ni(I1) > Pb(I1) > Co(I1) > Zn(I1) > Cd(I1). chlorate; lower gra h, K c for the ratio of the Cu(yF: EDTA to Ni(I1)-EDkA ions in potassium perchlorate.
TABLE I STABILITY CONSTANT RATIOSIN POTASSIUM NITRATE fi = 0.1; t = -30' Ratio
.
CuY/NiY NiY/PbY CuY/PbY NiY/ZnY
log
KO
0.315 f 0.02 0.71 f .02 0.88 =k .03 2.20 f .03
Ratio
log KO
CoY/ZnY 0.07 f0.01 CoY/CdY 0.30 f .07 NiY/CdY 2.46 =k .06 PbY/CoY 2
Effect of Ionic Strength.-The effects of decrease in ionic strength are given in graphic form in Figs. 3 and 4, In Fig, 3 it is aeen that the ratio of stad
bilities of copper(I1) to lead(I1) chelates decreases sharply as the ionic strength of the potassium nitrate medium decreases. There is a somewhat less rapid decrease of stability constant ratio in potassium perchlorate solution. Both curves seem to approach the same value in dilute solutions, indicating a thermodynamic ratio of 2.3 : 1 at infinite dilution. The nickel(I1)-lead(I1) ratios in Fig. 4 show little change over the range of ionic strengths studied in both media. At lower ionic strengths, however, it seems that the ratios in potassium ni-
DETERMINATION OF ETHYLENEDIAMIKETETRAACETATE CHELATION
Oct., 1953
trate probably break sharply and become equal to the potassium perchlorate value at infinite dilution. The copper(I1)-nickel(I1) ratio in perchlorate solution, also shown in Fig. 4, has a definite slope analogous to that of the copper(I1)-lead(I1) system. Absolute Stability Constants.-The results obtained for the stepwise dissociation of the nickel(I1) complex a t low pH are given graphically in Fig. 5. It is seen, that as the hydrogen ion concentration is increased the intensity of absorption is greatly decreased, but the shape of the absorption band remains the same. The absolute stability constant for the NiY complex was calculated by the method outlined above, and from this the absolute stability constants were evaluated. The results are given in Table 11. TABLE I1 METALCHELATE FORMATION CONSTANTS
Metal
log K
Ni
17.5
cu Pb co
Zn Cd
17.5 17.8 17.7 17.2 16.8 15.4 15.4 15.3 15.3 15.0 15.0
Remarks
Reference value Perchlorate sol. Nitrate sol. Perchlorate sol. Nitrate sol. Perchlorate sol. Nitrate sol. Reference value Perchlorate sol. Nitrate sol. Perchlorate sol. Nitrate sol. Perchlorate sol. Nitrate sol.
log K (Schwarsenbach) in 0.1 M chloride solution a t Z O O
18.45
697
% 52
1000
950 1.
Fig. 5.-(A) Extinction coefficient of Ki(I1) chelate a t pH 4.5. Formal extinction coefficients of Ni(I1) chelate in (B) 0.009 ill HC104, (C) 0.018 Ai HCIO, and (D) 0.025 df HClOc
TABLE I11 RELATIVE STABILITIES CALCULATED BY DIRECT AND INDIRECT METHODS log K (direct)
log K (indirect)
Cu Ni Ni Pb
0.63
0.55
f,-G;il
Cu/Pb
cu Ni
0.38
0.37
p =
16.1
“INi
e P b / P5 b
0.29
0.37
EG;i1
16.2
Cu/Ni
w-pb Ni
0.37
0.38
p
=0
16.5
Ni/Pb
pb/E cu cu
0.34
0.26
Ni/Pb
21%
0.27
0.26
p
=0
Ni/Cd
?!!.%/cO Zn Cd 5
2.45
2.43
p
=
0.18
0.31
$--$*’
18.34 18.2
The dissociation of the copper(I1) complex at low pH was also investigated. It may be seen that as the hydrogen ion concentration is increased, the curve shifts from its original shape. The maximum becomes broader at 700-760 mp, and above 900 mp the formal extinction coefficient rises above that of the normal complex. Apparently, therefore, one or more additional complex forms are present a t low pH values. These may be hydrogen complexes similar to that described by Schwarzenbach and Heller2 for the ferrous ethylenediaminetetraacetate complex.
Discussion As a check on the self-consistency of the stability constant ratios, a series of interrelation tests were made by calculation of one ratio from two others involving a common metal. Thus the copper(I1)-nickel(I1) ratio is obtained by dividing the copper(I1)-lead(I1) ratio by the nickel(I1)lead(I1) ratio. Results of such comparisons, given in Table 111, show that the relative stabilities calculated by both the direct and indirect methods agree quite well. These results indicate that the spectrophotometric method is a t least as sensitive as other methods previously employed for the determination of stabilities of transition metal-ethylenediaminetetraacetate chelatea.
Direct ratio
Cu/Pb
Cu/Ni
Indirect ratio - 1 -
-*-
Ni
Ph
&&
Conditions
0
KNOi
The sharp drop in the relative stabilities of the copper(I1) t o lead(I1) chelates in potassium nitrate as the ionic strength is decreased indicates the possibility of nitrate complexing. This idea is strengthened by the much lower slope for the same relative stabilities in potassium perchlorate solution, since the perchlorate ion is generally assumed not to form complexes with most bivalent metal ions. If consideration is restricted to the first nitrate complex of copper(I1) and lead(I1) ions, the corresponding formation constants KcU and Kpb may be substituted into equation (2) to give
+
where ZPb and ZCu represent [Pb+2] [PbN03+] and [ C U + ~ ] [CuNOa+], respectively. These are the values obtained experimentally for the free metal ion. If ib is assumed that there is no com-
+
VINCENTL. HUGHES AND
698
plexing by perchlorate ion, K’c can be evaluated from the KO value obtained at the same ionic strength in perchlorate solution. Hence, for the nitrate solutions (9)
This relationship may be rearranged to the linear form
Substitution of experimental values of KO,K’, and [Nos-] into this equation should give a series of straight lines intersecting at the common solutions for Kcu and Kpb. However, examination of the form of equation (9) indicates that the method is insensitive when the constants and nitrate concentration are small, since this would make the second term in both the numerator and denominator of the fraction small compared to unity. The equation is also insensitive when the nitrate complexing constants are nearly the same. Both sets of conditions apply in the present instance. Further, in order to solve (9) it was necessary to use data at very low nitrate concentration, The constants thus calculated varied by 50% and were several times the magnitude usually accepted for the metal ions investigated. On the other hand, the ratio of nitrate complexity constants determined from the data at 0.1 M nitrate concentration gives the ratio of nitrate complex formation constants as 6:3:2 for PbN03+:CuN03+:NiN03+. It is be-
l
N O
I
C”
2.
17
Fig. &-Chelate stability constants and ionization potential, as a function of atomic number. Solid lines: (A) log klk2 of acetylacetone, (B) log klkz of ethylenediamine, and (C) log k of ethylenediaminetetraacetic acid chelates. Dotted line: second ionization potential (E*).
ARTHUR E. MARTELL
VOl. 57
lieved that these relative values are fairly accurate. Although the measurement of the interaction of metals with nitrate ion is not the primary purpose of this research, it seems that evaluation of the absolute magnitude of the equilibrium constante would be possible by extension of the measurements to higher nitrate concentration. The values of the absolute stability constants reported in Table I1 for nitrate solution were calculated from the measured relative stabilities and the absolute constant for the lead(I1) nitrate complex calculated by the equation K’M = K M (1
+ K Y N O[Nos-] ~ )
where K‘M is the constant determined in perchlorate solution, K M is the constant determined in nitrate solution without correcting for nitrate comwas ~ assumed to be 15, the plexing, and K M N O value reported by Harned and Owen15 as the nitrate complex formation constant for lead(I1) at 25’ and 0.1 p . Thus, the values reported in Table I1 for nitrate solution are uncorrected values, which may be employed directly to calculate metal ion concentrations in 0.1 M nitrate solution. The perchlorate constants listed in Table I1 for Co(II), Zn(I1) and Co(I1) were calculated on the assumption that the ratios for these metals in perchlorate solution are the same as in nitrate solution. This assumption is justified by the form of equation (9) and the small nitrate complexity constants for these metals. The effect of ionic strength on the stability constant ratios in perchlorate solution must ultimately be resolved in terms of the activity coefficients of the species involved, and so it would be well to consider the similarities and dissimilarities of the corresponding molecular structures. In the case of nickel and lead, the activity coefficients of the two metal ions would be expected to show about the same dependence on ionic strength. Thus, the ratio should be close to unity. The two complexes are probably hexadentate, and thus completely enclosed by the complexing anion. Since the structures and the charges are so similar, we might expect the activity coefficient ratio to remain close to unity. The over-all effect would then be small, and this is apparently the case experimentally. The marked dependence on ionic strength of the copper (11)-lead (11) and copper(I1)-nickel(I1) systems is interesting in view of the fact that the metal ions involved are of the same valence type. A similar system,l6 copper(I1) and nickel(I1) sulfate, actually shows exactly the same dependence on ionic strength up to 0.7 molal. In each case one entity, the lead(I1) or the nickel(I1) complex, is hexadentate, while the copper(I1) complex is square planar. A structural difference of this type would be expected to give rise to appreciable differences in the activity coefficient, except in an infinitely dilute solution. Much of the discrepancy between stability constants obtained in this investigation and those reported by Schwarzenbach and Freitag’ is due to the (15) H. Harned and B. Owen, “The Physical Chemistry of Electrolytic Solutions,” second edition, Reinhold Publ. Corp., New York, N. Y., 1950, p. 147. (16) See reference 15, p. 427,
.
Oct., 1953 I
THESTRUCTURE OF OXYGEN FLUORIDE
699
difference in the value employed for the fourth ands show similar behavior with the metals in Lhis dissoclstion constant, kq, of ethylenediaminetetra- series in that the stabilities increase as the atomic acetic acid. The insertion of Schwarzenbach's number increases, with a maximum a t copper(II), value of IC4 into equation (7) brings our results into followed by a sharp drop in stability for the zinc(I1) The ethylenediaminetetraacetate better agreement with his: Ni, 18.0; Cu, 18.3; complexes. Pb, 17.7; Co, 15.9; Zn, 15.8; and Cd, 15.5. It chelate differs remarkably from the others shown in is to be expected that some differences in the re- Fig. 6 in that the plot is much flatter, and the results would remain, in view of the fact that the agent is correspondingly much less selective for the present investigation was carried out at a higher transition metals. The correlation of chelate stability constants of temperature and in a different electrolyte. The order of stabilities obtained in this investigation the transition metals with the second ionization also differs somewhat from that determined by potential of the metal, first suggested by Calvin Schwarzenbach and Freitag. However, these dis- and Melchior,18has been found to hold for a wide crepancies are not surprising, in view of the strong variety of ligands. The ionization potentials effect of a supposedly inert electrolyte on the values plotted in Fig. 6 as a function of atomic number, of the relative stabilities of the chelate systems indicate that this correlation holds for all the metals and ligands shown. The correlation with ethylillustrated in Figs. 3 and 4. It has been pointed out by Irving and Williams" enediaminetetraacetate chelate stabilities differs that the divalent metal ions exhibit the same order from the others shown in that the properties of of stability with a large number of ligands, e.g., the metal ions, as measured by the ionization ethylenediamine, propylenediamine, ammonia and 'potentials, apparently have much less influence salicylaldehyde. This order of stability, Cu > Ni on chelate stability than is the case for the other > Co, Zn > Cd, is the same as that obtained in the ligands. present investigation. Acknowledgment.-The authors are indebted to The stabilities of various chelates of a number of the U. S. Navy Office of Naval Research for the the first row transition metals are plotted vs. atomic support of this work under contract Nonr-596(00). number in Fig. 6. It is apparent that all the lig(18) M. Calvin and N. C. Melchior, J . Am. Chem. Soc., 7 0 , 3270 (17) H. Irving and R. Williams, Nature, 162, 746 (1948).
(1948).
THE STRUCTURE OF OXYGEN FLUORIDE BY JAMESA. IBERS' AND VERNER SCHOMAKER Contribution No. I774 from the Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, California Received Januarv B6, 1068
A bond angle of 103.8 f 1.5' and a bond length of 1.413 f 0,019 A. have been found for oxygen fluoride by the method of electron diffraction. Combining our data with the spectroscop$ rotational constant reported by Bernstein and Powling,2 we arrive at a bond angle of 103.2' and a bond length of 1.418 A. as most probable values. Special attention was given the problem of weighting the measured ring diameters, and a simple adjustment was finally adopted which may prove useful in other electron diffraction work of this kind.
The bond length and bond angle in oxygen fluoride (OFz) are of interest., especially in relation t o the values for other simple compounds of the electronegative elements, but they have not been precisely determined. The best previous values are dependent on early electron diffraction results and on a single spectroscopically determined molecular constan t . For the present electron diffraction reinvestigation Professor L. Reed Brantley of Occidental College kindly gave us a one-to-one OF2-02 mixture prepared by the usual method.8 After nearly all the oxygen had been pumped off' at liquid-nitrogen temperature,4 the sample was fractionally distilled at (1) National Science Foundation Predoctoral Fellow, 1952-1953. (2) H. J. Bernstein and J. Powling, J . Chem. Phys., 18, 685 (1950). (3) P. Lebeau and A. Damiens, Compt. rend., 188, 1253 (1929); 0. Ruff and Menzel, 1.anoru. u. allgem. Chem., 190,257 (1930); G . H. Cady, J . Am. Chem. Soc., 37, 246 (1935). (4) The fact that OF1 can be obtained in about 95% purity bypump-
W.
ing off the oxygen a t liquid-nitrogen temperature has been confirmed by the recent work of J. G. Schnidein. J. L. Sheard. R. C. Toole and T. D. O'Brien, THIS JOURNAL. 56,233 (1952).
about - 160" through a 65-cm. vacuum-jacketed column; iodometric analyses of the product ranged from 98.5 to 99.4 mole % OFz. Photographs were made and interpreted in the usual way,5 except that in view of the simplicity of the problem the visual and radial distribution curves were not drawn. The camera distance and elgctron wave length were 10.91 em. and 0.06056 A.! respectively. Intensity curves were calculated for rigid symand metrical OF2 models with 0-Fo = 1.42 F . F ranging from 2.10 to 2.32 A., for various admixtures of 0 2 . Figure 1 shows three of the curves for pure OFz with critical markse to indicate important comparisons with the photographs. The more sensitive features, important for the angle determination, are marked on curves A and C, and the less sensitive features on curve B. Finally, the measured ring positions were compared with
i.
+
(5) K. Hedberg and A. J. Stosick, J . A m . Chem. rSoc., 74,954 (1952). (6) W. F. Sheehan, Jr., and V. Schomaker, ibid., 74,4468 (1952).
