The large probable errors in A H o ( a ) and ASo(a) are not unreasonable because slight departure from equilibrium would cause the mole fraction of Ice&& to be in error by relatively large percentages in such a dilute solid solution region. The solid curve in Fig. 1 shows the relationship between oxygen dissociation pressure and temperature when a hematite phase is in equilibrium with a magnetite phase. The corresponding data of Sortong and Darken and Gurry5 are also shown. When values of Po, and T from this curve are substituted along with the above values of AHo(cu) and ASo(a)into equation 15, the composition (ie., X1 and X 2 ) of the hematite phase in equilibrium with the magnetite phase can be calculated a t temperature 1'. This has been done and the results are plotted in Fig. 3. It is evident from the curve in Fig. 3 that the solubility of Fe304 in the hematite phase is very slight. G. Conclusions.-The pressure-compositiontemperature relationships a t high temperature for the wiistite, magnetite and hematite phases of the iron oxide system have been successfully explained by consideration of the entropy of mixing of lattice defects in the lattice. I n this study the (9) F. J. Norton General Electric Reaearch Laboratory Report No. 55RL-1248.
heat of mixing of the defects with the lattice has been assumed to be zero. The experimental results agree with this assumption. Another helpful assumption has been that the charge distribution in a given phase such as wiistite is independent of the composition of that phase. This has been a valuable guide in postulating horn the defects are positioned in the lattice. By knowing the expression for entropy of mixing of defects as a function of oxygen content, it is possible, by applying thermodynamic principles to the experimental data on oxygen dissociation pressure, composition and temperature, to obtain standard heats and entropies of oxygen dissociation from the wiistite, magnetite and hematite phases. By use of the values for the standard heat and entropy of oxygen dissociation and the corresponding expression for entropy of mixing of defects as a function of composition, one can readily calculate the oxygen dissociation pressure for any oxygen content and temperature of wiistite, magnetite and hematite phases. Acknowledgment.-The author gratefully acknowledges the support of this work by the u. S. Army Signal Corps Engineering Laboratories, Fort Monmouth, New Jersey, under SC Cont'ract NO.DA-36-039-SC-74904,
KINETICS OF THE REACTION BETWEEN HYDROGEN PEROXIDE ASD AQUO-(ETEIYLENEDIAMINETETRAACET0)-COBALT(I1) AND EVIDENCE FOR THE FORMATION OF A PEROXODICOBALT(III,III) COJIPLEX' BY RICHARD G. YALMAN Department of C h i s t r y , Antioch College, Yellow Springa, Ohio Received November 8,1960
A psroxodicobalt(II1,III) intermediate is formed during the oxidation of aquo-(ethylenediaminetetraacet0)-cobalt(11) t o cobalt(II1) by hydrogen peroxide. In the presence of a large excess of cobalt(I1) the oxidation is stoichiometric. Otherwise oxygen is formed. The kinetics of these reactions were determined by manometric, spectrophotometric and polarographic techniques. The mechanisms of these reactions are discussed.
oxidation of cobalt(I1j to cobalt(II1) is independent of the product(sj at this wave length.
(1) Part o f the material in this paper was presented before the Division of Physical Chemistry at t h e 134th meeting of the Amerioen Chemical Society, Chicago, 1958. (2) G . Sohwarzenhaoh, Helv. Chim. Acta, 83, S O (1040). (3) I. A . Shimi ar.d W. C. E. Hipginson, J . Chem. Soc., 260 (1958). (4) S. M. Jorgonsen. Acta Chem. Scand.. 9 , 13G2 (1965).
Reagents.-Dihydrogen aqno (ethylenediaminetetraaceto)-cobalt(I1) was prepared by the method of Shimi and H i g g i n ~ o n . ~This complex was the only source of cobnlt(I1) and ethylenediaminetetraacetic acid wed in these experiments. Crystalline catalase was obtained from hrutritionnl Biochemicals Corp., Cleveland, Ohio. All other chemicals were C.P., A.C.S. reagent grade. All salt Eolutions were carefully filtered and, where necessary, analyzed by standard volumetric procedures. Reaction mixtures were prepared by rapidly mixing solutions of the cobalt(I1) complex with solutions of hydrogen peroxide. The latter contained sufficient phosphate buffer and sodium perchlorate so that on dilution t o volume each reaction mixture was 0.1 M in phosphate and had a total ionic strength of 0.5 M . In a number of experiments excess hydrogen peroxide was destroyed after 5 minutes by chilling the reaction mixture and adding catalase. Solutions treated in this way will be called catalase treated reaction mixtures. In one series of experiments t o determine stoichiometry dilute solutions of hydrogen peroxide were added dropwise with continuous stirring to buffered solutions of the cobalt (11) complex. The pH's of ail solutions were measured periodically with a Beckman Model G pH meter.
When hydrogen peroxide is added t,o solutions aquo- (et liylenediaminetet raacet,o)-cobalt,(11), YCOH~O-~, cobalt(II1) complexes are formed.2 At the same time hydrogen peroxide is decomposed. The purpose of this investigation was to examine the kinetics of these reactions. Preliminary experiments showed that the evolution of oxygen could be studied by manomet,ric techniques in the pH range of 6.5-8.5 a t 30'. Under these conditions there is an equilibrium between aquo-(ethy1enediaminet'etraaceto)-cobalt(111), YCoHZO-, and hydroxo-(ethylenediaminetetraaceto)-cobsalt(III), YCoOH-,* and both of these complexes form et'hylenediaminetetraacetocobalt(III), Y C O - . ~ ~At ~ 576 mp3,4these complexes have t,he same molecular absorption coefficient and t,he spectrophotometric study of the
HYDROGEN PEROXIDE ASD AQUO-(ETHE-LEXEDIAMIXETETR.L~CETO)-COBALT(~~) 557
Oxygen Experiments .-The amount and rate of oxygen evolut,ion was determined with an Acme Equipment Co. Warburg Apparatus using Summerson manometers and onearm Warburg flasks. The latter were calibrated by the r e a d o n between hydrogen peroxide and potassium permanganate in dilute sulfuric acid. The reaction mixture was generated by dumping solutions of the cobalt(I1) complex from the side arm into buffered hydrogen peroxide solutions. The thermobars conkained similar solutions, but no cobalt(I1). After a short induction period straight line curves were obtained by plotting log ( H , - H t ) against t , where H , is the total cliange in the height of the manomet.er. First-order rate constants were calculated from the slopes of t.hese lines. In a few experiments catalase treated reaction mixtures were placed in Warburg flasks and, after equilibrating for 15 minutes, the evolution of oxygen was observed. Although the amount of oxygen formed in these experiments was only about 40%, of t,hat formed in the original reaction mixtures, the rate of oxygen evolution was nearly twice as great, indicating a reaction involving catalase. In duplicate experiments the same amount of oxygen was formed almost instantaneously when potassium triiodide was added from the side arm t,o catala& treated reaction mixtures. Spectrophotometric Experiments.-Optical density measurements were made with a Heckman Model DU Spectrophotometer with a phot,omultiplier attachment using 1 cm. silica cells. The cell compartment was maintained a t 29.7 + 0.1" by circulating water from a constant temperature bath throilgh coils in thc cell housing. A number of measurements were made a t 0.5" by circulating water from an icebath through the cell housing. In these experiments Drierite RBR placed in the cell houeing to prevent fogging. The absorpt,ion spectra of YCOOH-~,YCoH20-and YCowere determined from 320 to 660 mp. The results were in agreement with those reported in the l i t e r a t ~ r e . ~The .~ concentration of cobalt( 111)was determined from measurements a t 5T6 mp, where the complexes have the same molar :tbsorption coefficient. In the experiments to determine stoichiometry corrections were made for the amount of unreacted cohalt(I1). ,4t 576 mp thc optical density of the reaction mixtures increases raaidlv and then more slowlv. From d o t s of log (D, - &)