Dec., 1961
KISETICSOF SISTERING OF PLATIXUM SUPPORTED ON ALUMINA
The process for the formation of the ion of m/e = 146, Sn(CH&+, from tetramethyltin, is considered to be the same as that for the lead and silicon compounds, with two neutral CH3 fragments formed. Consideration of the approximate energy necessary to remove a methyl group, and of the fact that the energy for t h e accompanying formation of an ethane molecule is very great, leads us to conclude that the neutral fragments are two CH3 radicals. Thus, AHf+(SnlVIe2)= 225 kcal./mole. The heat of formation of Pb(CH&+ is calculated to be 236 kcal./mole. Energetics indicate that the process in Table 111, as given for this ion, is the dominant one, invohing two methyl groups as neutral products of the dissociative ionization process. M(CH3)3+.--'I'hese ions are the dominant species in the mass spectra of all of the compounds considered in this E,tudy. The ion of m/e = 7 3 from tetramethylsilicon is SiC3H9+. We presume that the structure is that given by Si(CH3j3+.Energetics dictate that th12 neutral fragment is the methyl radical. On this basis, we calculate AHf+(SiMe3) = 166 kcal./mole. The heat of formation for the dominant species in the tetramethyltin spectrum, Sn(CH3)3+,is 183 kcal./mole for the process involving the formation of the neutral fragment CHB. From energetic considerations, Ihis is the only possible process for the formation of this ion. Pb(CH3)3+is formed by the removal of a methyl group subsequent to ionization of the parent molecule. The heat of formation for this ion is calculated to be 206 kcal./mole. M(CH3)4+.--'.rhis is the parent molecule-ion in the compounds studied. In each of the spectra, this ion is very low in abundance. The heat of formation calculated for Si(CH3),+ from the determined appearance potential is 163 kcal./mole. The heat of formation for Sn(CHa)4+is177 kcal./ mole; the process of its formation involves the removal of a single electron. Because the SnlZ1 isotope does not occur naturally, the appearance potential of the ion with m/e = 180 definitely is to
2159
be attributed to Sn1ao(CH3)4+and not to Sn12i(CH3)3(CH2)+. The ionization potential of Sn(CH3)4was found to be 8.25 f 0.15 e.v. The parent ion of tetramethyllead, Pb(CH3)?+, is of extremely small abundance, and the ionization potential was quite difficult to determine. Only fair agreement was obtained between numerous determinations; therefore an error range of k0.4 e.v. was assigned to the value of the ionization potential determined. The agreement of this experimental result with the calculations made on the basis of the modified equivalent orbital treatment substantiates our experimental result (see Table IV). Thus, AHf+(PbMe4) = 217 kcal./ mole. The ionization potential of 8.0 f 0.4 e.v. herein reported is 3.5 e.v. lower than that reported by Fraser and Jewitt.8 The values reported for the ionization potentials of a number of the biscyclopentadienyl-me tal compounds 2 2 suggests that the ionization potential for tetramethyllead is indeed much lower than 11.5 e.v., and should be closer to that of the gaseous metal atom. Again, our result of 8.0 f 0.4 e.v. agrees with this conclusion. It may be noted that the ionization potentials of the tetramethyl organometallics of Group IV decrease in a regular fashion, as expected. We hope to be able to study experimentally tetramethylgermanium in the near future; it is anticipated that the fragmentation processes will be the same as those observed in this study. From interpolation of the heats of formation of the various ions and by back-calculation, as well as from the interpolation of the appearance potentials, it is estimated that the appearance potentials will be approximately 19, 16, 13.5, 10.5 and 9 e.v., respectively, for the formation of Ge+, GeCH3+, Ge(CH&+, Ge(CH3)3+and Ge(CH3)4+from tetramethylgermanium. Acknowledgments.-The authors wish to thank Dr. R. L. Soulen for his aid in the preparation of the samples of tetramethyllead and tetramethyltin. ,41so, we wish to thank Dr. J. L. Franklin for his helpful comments and criticisms of this study. (22) L Friedman, A . P. Irsa and G. Wilkinson, J . .4m Chem. Soc., 77, 3689 (1955).
THE KINETICS OF SINTERING OF PLATINUM SUPPORTED O S ALU31IX;A BYR. A. HERRMANX, S. F. ADLER,M. S. GOLDSTEIN AND R. M. DEBAUX American Cyanamid Company, Stamford, Connecticut Receieed June d f i , 1961
Various samples (of Pt-AL03-Cl re-forming catalyst were subjected to a series of heat treatments Examination of the fresh and heated mctterials by measurement of chemisorptive capacity for H?,rate of chemisorption of Ho,P t crystallite size by X-ray line broadening and measurement of Pt solubility in HF lead to tho conchsion that Pt is present in the original catalyst in a highly dispersed, perhaps ionic, form and that heat treatment causes the formation of Pt crystallites, and not merely their growth.
Introduction The state of djspersion of platinum in platinumalumina re-forming catalysts, both impregnated and cogelled types, has been the subject of an earlier inl'estigation at this Laboratory.' fb'drogen (1)
s. F.
Adler anrl .I J. K e a i n e y , .I.
w A y 3 .Cjiern
, 64,208 ( 1 ~ 0 ) .
chemisorption measurements and X-ray diff raction analyses were employed to describe the nature of platinum and the effect of various treatments on it. These included thermal and hydrothermal treatment R, exposure to high temperature hydrogen--hydrocarbon atmosphere (re-forming),
2190
R. A. HERRMANN, S. F. ADLER,M. S. GOLDSTEIN AND R. M. DEBAUN
Vol. 63
Catalyst €3 always was run simultaneously with sample A The samples were intimately mixed before placing them into the quartz tube, thus ensuring as precise a comparison as possible under exactly the same conditions of temperature and atmosphere. Subsequent separation of these catalysts was greatly facilitated by their different geometrical forms (cylinders and beads), which allowed mechanical separation. Chemisorptive Measurements.-The procedure and ap paratus used for measuring hydrogen chemisorption on supported platinum catalysts already have been described in some detail.’ Briefly, the chemisorption of hydrogen was measured volumetrically at a pressure of about 9 mm. which was maintained constant to within approximately 10%. The measurements were made with the sample a t 200 . Sample preparation consisted of a preliminary outgassing for two hrs., then a short (20 sec.) pretreatment with hydrogen at 500°, followed by overnight,?umping to a final pressure of approximately 10-6 nim. 1he uptake of hydrogen a t 200’ then was followed with time and the volume of hydrogen (STP) adsorbed per gram of platinum at two hr. exposure was taken as a measure of the chemisorptive capacity. X-Ray Diffraction Line Broadening.--A standard Philips diffractometer was equipped with a “tracergraph” print timer and a programmer to allow point-by-point intensity measurements, recording the time required to accumulate a preset number of counts at a scattering angle, and automatically shifting the diffractometer to the next scattering angle. The reciprocal of the time measured was plotted as a function of scattering angle over a range to cover the platinum (311) reflection. Unfiltered copper radiation and largest slits available were used throughout to ensure maximum intensity from this, the third strongest platinum reflection. Platinum crystallite size determinations %-eremade by standard line broadening techniques using the plotted data. Solubility of Pt Species in HF.-One to 2 g. samples of Preparation of Catalysts.-All Type A catalysts were pre- powdered catalysts were treated with an escess of an appared in a fashion similar to that previously described’ under proximately 25% aqueous solution of HF according to an “Impregnated” Catalysts. Subsequent to the normal prep- established procedure.5 The mixtures were filtered after a t aration, samples AB and A4 %-ereprepared from A1 and -q2 least one hr. and the filtrates were analyzed for Pt spectroby washing with an ammonia solution to remove chloride, photometrically. thus modifying the Pt/C1 ratio. Catalyst B is a commercially used re.forming catalyst. Results The compositions of this material and all Type A catalysts Table I1 gives the evaluation of two of the cataare given in Table I.
and ordinary catalytic regeneration procedures. The work was independently corroborated,2 even to quantitative aspects. In these and other paper^^.^ the picture shown, for impregnated catalyst, is of a highly dispersed form of platinum, which undergoes transformation to a clustered and “crystalline” form upon prolonged exposure to various atmospheres a t high temperature. It has been of interest to study the effect of heat treatment on platinum re-forming catalysts of commercial types for two reasons. First, a study of the kinetics of observable changes on heat treatment and the fitting of these into a plausible mechanism may well enable more subtle conclusions to be drawn as to the nature of the reacting material (supported platinum) as originally obtained and as governed by various preparation techniques. Moreover, the stability of the platinum promoter is related to catalyst inactivation in use and is therefore of considerable practical significance. Our samples of platinum catalysts therefore were subjected to various heat treatments and examined by X-ray line broadening and by hydrogen chemisorption techniques. In the latter technique, both the absolute amounts sorbed and the rates of sorption were studied. Experimental
TABLE I CATALYST COMPOSITIONS % Platinum
% Chloride % Fluoride Surfacearea,m.*/g. Pore volume, cc./g. Alumina
Form
A Ai Ai AI A4 B 0.774 0.787 0.315 0.720 0.300 0.375 0.830 0.920 0.480 0.110 0.040 0.350 .. .. 0.350 225 199 202 199 202 176 0.617 0.625 0.617 0.625 0.940
... Y
.
...
Y
1/16”D X 3/16’L
Y
... Y
cylinders
.
Y
Y
1/16@beads
Heat-Treating Procedure.-Unless otherwise noted, all samples were subjected to a standard calcining procedure of heating to 593’ for one hr. in air before division into small samples for further heat-treating and/or chemisorption measurements. Samples to be heat-treated were placed in a quartz “U” tube through which a steady stream of prepurifiFd nitrogen was continuously passed a t a rate of 200 cc./ min. Approximately 8 g. of catalyst was heat-treated a t a time except when samples A and B were treated simultaneously, in which case about 8 g. of each was used. After placing the sample in the “U” tube containing a thermocouple, the nitrogen stream was allowed to pass over the sample for at least 20 min. before insertion into the hot furnace. This ensured against excess moisture being present which would tend to steam the catalyst in the early stages of heating. The sample temperature was checked several times during a run by means of the thermocouplc buried in the catalyst. The furnace temperature was maintained constant to within ~ k 2 Owith a Wheelco stepless proportioning controller. Following the heat-treatment, samples were cooled in the quartz tube under flowing nitrogen. (2) L. Spenadel and M. Boudart, J . Phys. Chern., 64.204 (1960). (3) G. A. Mills, 9. Weller and E. B. Corneliua, Paper 113, Proceedings 2nd Intern. Catalysis Conference, Paris, France, 1960. (4) R. C. Pitkethly and A . G. Goble, Paper 91, ref. 3.
lyst samples after various heat treatments ur,der dry conditions. Each sample was examined by hydrogen chemisorption (after preliminary pretreatment with hydrogen) and X-ray diffraction. The rate of decrease in chemisorptive capacity was established to be second-order with respect to remaining chemisorptive capacity. This vias done using primarily the data for samples A and B, especially at 625’ where the data are more complete. Thus, having established the rate dependency, it was assumed to hold for samples AI, Az, A3 and A4 where there is insufficient data to make a judgment. Each datum was used together with the chemisorptive value of the initial (untreated) sample to calculate an apparent specific second-order rate constant for each material. This constant therefore has dimensions of K = (cc. HZ/g. Pt)-’ hr.-l. The consistency of the rate constant for sample A a t each of the various temperatures is evidence of the second-order relationship, allowing that two of the twelve points represent sufficiently wide deviations from the group to be discounted as incorrect. It shoald be noted that these data cover a range of loss of chemisorptive power in excess of 80% of the original (untreated) value, showing the observed kinetics to hold over a considerable degree of degradation. In addition to the total chemisorption, measured (5) K.W.McHenry, R. J. Bertolacini, H. 31. Brennan J. L. Wilson and 11. S. Scelig, Paper 117, Proc. 2nd Intern. Cata1,sis Conference, Paris, I’rance. 1960.
KINETICSOF SINTERING OF PLATINUM SUPPORTED ON ALUMINA
Dec., 1961
2191
TABLE I1 EVALUATION OF HFATEDCATALYSTS A AND B HI
Sample
Catalyst and treatment Ilr.
T, OC.
chemisorr tion, cc. Ha/& catalyst a t 120 min.
Chemisorption -parameters-a1
ai
A
Fresh 0.683 56.8 A Fresh .674 57 5 A 44 .431 69.6 564 A 47.5 233 .119 564 564 A 70.5 84.1 .372 .292 96 A 187 564 ,116 564 A 353 230 87.8 A 34 594 .326 134 ,215 594 A 48 594 .135 A 93 162 A 4 88.0 ,355 625 .141 625 A 8 190 625 ,123 192 A I. 8 625 .080 d 40 280 B Fresh 0.298 143 13 88.2 Fresh .303 B 44 564 .136 191 564 .057 B 47.5 500 564 328 B 70.5 .045 .051 564 411 B 167 .030 564 562 B 353 B 24 .117 594 161 264 .017 594 B 48 .043 B (33 594 460 B 4 625 .088 237 .035 625 585 8 B 625 .030 B 18 418 B 40 625 .023 574 0 By X-ray line broadening. Assumed to be incorrect.
*
after 120 minutes of exposure to Hz, the rate of uptake of hydrogen in the chemisorption test also was followed. In the calculation of the chemisorption kinetics, the q vs. t data were fitted to equation 1.6 Accordingly, q values were plotted against log t in minutes, as illustrated in Fig. 1. From the slopes of these plots, generally almost linear, values of cy1 and al were estimated from the values of q a t 10 and 100 minutes.
- (log t + log a1 LYI) assumed alcyltl > 1, as indeed p :=
2.3 ff1
This model it was found to be. As can be seen from Table 11, the chemisorptive capacity of both samples A and B decreases markedly with both temperature and time of heating. A regularity in this decrease is noted in the opportunity of fitting a second-order model to this data. I n Figure 2, the obtained rate constants are plotted against reciprocal absolute temperature to obtain activation energies of about 70 kcal./mole. It is interesting to note that both samples A and B show similar behavior as regards decrease with time of chemisorptive power, even to the effect of temperature. It is noted that examination of sample AI (a nominal dup1icai;e preparation of A) , gave similar results, as shown in Table 111. Also shown in this table are the results with a version of A containing ( 6 ) M.
J. D. Low, Chem. Revs.,
60, 267, (1960).
x
10s
71 68 33 6.8 25 7. 6 5.5 21.3 11.3 7.9 22.7 6.9 6.5 4.1 13 21 6.8 2.2 3.2 2.7 1.9 7.6 6.0 2.3 4.9 1.8 2.5 1.8
Pt crystallites" Relative Sire, A. amount
...
...
nil nil
Second-order Rintering rate constant 104 (CC. Hz/g. Pt) -1 hr. - 1
...
...
160
2.9
1.5 11.3' 1.3 0.9 1.6 5.2 5.1 5.0 26.0 55.0' 25.0 21.0
...
id
...
170 180 140 165 180 150 150
...
0.7 1.5 2.3 1.0 1.7 2.1 1.0
nil
180 170 150 155 165 130 145
0.4 0.7 1.0 0.7 0.8 0.7 0 7
130
1.1
... 3.5 11.4 10.0 3.8 3.2 8.3 7.5 8.2 77.0 120.0 63.0 38.9
30.43 8 0.42
& 0.41 .0.40
$ 0.39 0.38
3 0.37
9 0.36 0.14
$ 0.13
60.12 n" 0.11
1
2
Fig. 1.-Sample
10 20 100 200 Exposure to Ha, min. plots of chemisorption data. Samples pretreated at 44 hr. at 564'.
less platinum (Az) and also similar catalysts having much lower halogen-platinum ratios. Similar values for both the initial specific chemisorption and also the decrease in chemisorption were observed, although the initial specific chemisorption was decreased somewhat due to the halogen removal procedure. Table I1 indicates that decrease in Hz chemisorption is accompanied by the appearance of platinum crystallites and that the approximate amounts of crystalline platinum increase with increases in the severity of thermal treatment. It also is interesting to note that this regime does not appear to in-
R. A. HESRMANN, S. F. AD LEE^, M. S. GOLDSTE~N ABD R. M, DEBAUN
2192
Vol. 65
TABLDI11 EVALUATION OF OTHWRHCATTREATED CATALYST SAMPLES Catalyst
-Pt
--TreatmentHr.
“C.
fresh 627 627 frcsh 625 625 625 fresh 626 625 fresh 625 625
A1 Ai -4I
4 16
At
Az Az
6 12 28
123
A3 A2
6 20
Ah A4
A-4
5 20
&we,A.
CryystalhtesRelative amount
180 180
1.4 2.3
...
...
nil .7 .5 .9 1.6 2.1 nil nil
160
.4
150 200 195 175 175
...
H2 Chemisorption, cc. H l / & catalyst a t 120 min.
0.888,0.839 0.344 0.152 0.295,0.269 0.165 0.115 0.107 0.390,0.396 0.200 0.094 0.159,0.171 0.063 0.666
Sintering rate constant X 10‘ (cc. H d g . pt1-1 hF.+
... 34 27
... 13 14 6.5
... 35 29
... 59 16
genated catalysts (A3 and h4)possess lower hydrogen/platinum ratios (near one atom H/atom Pt) they are still more highly dispersed as regards platinum than typical cogelled catalysts. The regularity in decrease of the chemisorptive capacity with time is related to a second-order mechanism with respect to remaining chemisorptive capacity. In the event that the sintering of the platinum is ascribed to a coalition of small crystallites of platinum or to a growth of such crystallites by addition of platinum atoms from an “amorphous” phase, then such second-order dependence on the number of particles would reflect a growth rate depending on the sixth power of the surface I-C 0 .area of the crystallites and the sixth power of the \ chemisorptive capacity. This result is contrary to the observed second-order dependenae. What is possibly a more attractive explanation is one in \o which the chemisorptive capacity of platinum a0 - - t crystallites is taken to be low relative to that of the dispersed platinum and where the original disE persed platinum bodies are sufficiently small so $ 5 that the ehemisorptive capacity is primarily a 1.12 1.14 1.16 1.18 1.20 1.22 mewwe of their number since the crystallites repreI ~ K x. 103. sent a small percentage of the total surface area. Fig. 2.-Arrhenius plot of sintering rate constanta In such an event, the second-order decay of chemcrease the size of the crystallites beyond 150-200 A. isorptive capacity with time is explainable as indiAlthough there may be changes in the distributian cating a second-order dependence of the rate of reof crystallite sizes below this value, it would be dif- moval of platinum-hearing particles from a highly ficult to detect them by this method as the X-ray dilute “solution” or “dispersed” phase ta a “crystalaverage crystalhte size is a volume average crystal- line” or “condensed” phase. This notion is euplite size. parted by the fact that platinum crystallites in the The data also indicate pronouaced changes in form of three-dimensional aggregates which are the apparent values of the parameters in the chem- 150-200 A. in size will probably have a maximum isorption rate expression. These are reminiscent H/Pt ratio of about 0.14.2, thus contributing rery of changes noted with changes in the “cleanliness” little to the specific chemisorptive capacity of the of the adsorbing surface or the nature of the ad- material. We thus may think of a tmo-dimensorbing sites.6 sional “solution” of extremely highly dispersed Pt bodies (perhaps m atoms, ionic platinum, or groups Discussion Fresh catalyst samples A, AI, A3 and B all shaw of a few atoms) on the surface of the alumina, aggreH / P t chemisorptive powers of about one and one- gating b y a second-order law to produce deposits of half atoms of hydrogen per atom of platinum. microcrystalline platinum where the crystallites This indicates that the platinum is probably have dimensions of the order 150-200 A., each thus highly dispersed in the sample, perhaps even as containing 104-105platinum atoms. On examining the rate of chemisorption of hydro“monolayer patches” of platinum. This result is borne out by the absence of crystallites detectable gen as a funetioii of the thermal pretreatment of the by X-ray in the materials. Although the dehalo- catalyst, a regularity in the parameters of the rate i
:
\
i
\
\
Dec., 1961
KISETICSOF S I X T E R I N G O F P L A T I X U M
and 4 indicate that the parameters a1 and al of (1) appear to change quite regularly with changes in total chemisorptive capacity, p, due to various thermal pretreatments. In the general theory6 of the kinetics of chemisorption of gases on solids, the coefficient a1 is
"SOLUBLE"
TABLEIV RE-FORXIYG CATALYSTS
PLATINCM I N
A
593" for 1hr., subjected to Hz chemisorption A 593' for 1 hr. 625' for 4 hr. -4 593" for 1 hr. 625" for 40.5 hr. A 593' for 1 hr. 62.5' for 40.5 hr. subjected to H? chemisorption
83,80
2193
S U P P O R T E D ON h U M 1 N - k
-
5001 -
-
2
/
-
d
I
I
t
I I
/ I l l
chloride content by a factor of 20. The large variation in platinum to chloride ratio is shown to have little effect on the sintering rate. One then can infer that a chloroplatinate ion is not the only + 27 platinum species active in the sintering process. + The original presence of halogen may, however, + play some part in the distribution of platinum 23, 20 on the alumina surface during impregnation. The second-order rate constants are reported in While the crystallite size determination does not appear to add to the picture presented here, it in units of (cc. H2/g. Pt)-l hr.-l, which is related to no way contradicts it. Indeed, just as the same the fraction of the original chemisorptive capacity sort of mechaihm of sintering appears to be fol- remaining. If the distribution of the platinum lowed by catalysts A and B as judged by the time species on alumina were essentially uniform, then decay of chemisorptive capacity, so too the in- one would expect the kinetics of the sintering proccrease of crystalline X-rajr-visible material in both ess to be linked to the bulk concentration of platthe samples appears to have similar dependence inum. I n this case, k in units of (cc. Hz/g. catalyst)-I hr.-l, rather than in (cc. Hz/g. Pt)-' hr.-l, on both time and temperature. Examination of the rate constants which char- should be constant at different platinum levels. acterize the type A catalyst samples (Table 111) This would be so since the mean distance between shows little, if any, difference in sintering rates the fundamental platinum units mould decrease for these samples even though their platinum con- continually as their concentration incrwses in the tents vary by a factor of as much as 2.5 and the manner of the mean-free path of gas molecules with
+
56
2194
Vol. 65
A. V. CELIANO, M. CEPOLA AND P. S. GENTILE
increasing pressure. That this is not so can only be interpreted as meaning that the platinum is nonuniformly distributed and, in fact, must exist in local groupings of more than average density. Moreover, it has been shown that the number of adsorbing sites relative to total platinum (H/Pt) is independent of platinum concentration. This indicates that the nature of the local grouping of Pt species is unchanged with changes in concentration. As a matter of fact, the only change that could logically occur, in the light of the available data, is in the total number of such groupings. That is to say, increasing the platinum content of the catalyst appears to increase the number of "cities" of relatively fixed dimensions and Hz chemisorption characteristics.
A non-homogeneous distribution of platinum clusters on an amorphous alumina substrate is not unexpected, since it is well known that such a surface is heterogeneous from a crystallographic point of view. This is, in turn, associated with heterogeneity in the local surface density of hydroxyl groups and perhaps even with the relative acidity or basicity of those g r o ~ p s . ~ J Acknowledgments.-We wish to thank Mr. George Yates, who performed the heat treatments and hydrogen adsorption measurements, and Mr. William Doughman, who provided the X-ray data. (7) D. S. Rea and R. Lindquist, Paper 53, 136th Xational Meeting. Am. Chem. SOC., Atlantic City, September, 1959. (8) J. B. Peri and R. B. Hannan, J. Phvs. Chem., 64, 1526 (1960).
CHEMISTRY OF COGRDINATION COMPOUNDS. I. THE KINETICS OF FORMATIOIC' OF MONOACETYLACETONATOCOPPER(I1) ION BYALFREDV, CELIANO, Department of Chemistry, Seton Hall University, Souih Orange, A'. J .
MICHAEL CEFOLAAND PHILIP S. GENTILE Department of Chemistry, Fordham University, Bronx 67, AT. Y . Receioed June 67, 1961
+
The kinetics of the reaction copper(I1) ion acetylacetone F? monoacetylacetonatocopper( 11) ion in methanol wm investigated conductometrically in the temperature range -27 to 0'. The reaction obeyed second-order kinetics, first order with respect to copper(I1) ion, and first order with respect to the enol form of acetylacetone. The dependence of rate on hydrogen ion concentration indicates that the reaction does not involve prior ionization of the p-diketone, but rather the direct combination of the two reacting species. The enol content of the ligand as well m the energy and entropy of activation were evaluated.
Introduction The reactions between metal ions and P-diketones have been the subject of considerable investigation. I n the sphere of kinetics of complex ions interest has been focused mainly on the mechanism of substitution reactions rather than the rate of complex formation. Rates of reaction of various metals with ~-phenanthroline'-~and a,a'bipyridy14+ have been determined. Cook7 has studied the rate of complex formation of metal ions with aqueous 2-thenoyltrifluoroacetone and concluded that, for most metal ions investigated, including Cu2+, the rate of complex formation was identical with the rate of enolization. These investigators also studied the rate of complex formation using acetylacetone and 2,4-hexanedione and found that the rate of complex formation was not the same as the rate of enolization of these pdiketones. According to the principles enunciated by Bjerrum,s complex formation takes place in a series of reversible steps. J. A m . Chem. Soc., 78, 4211 (1956). (2) W. Brandt and D. Gullstiom, ibid., 74, 3532 (1952). (3) T. Leo, I. Kolthoff and D. Leussing, ibid.,73,3596 (1948). (4) P. Krumholz, J. Phys. Chem., 60,87 (1956). (5) J. H. Baxendale and P. George, Nature, 162, 777 (1948). (6) J. H. Baxendale and P. George, Trans. Faraday SOC..46, 736
M+n
+ HA + HA r'
MA+("-l)
+ I{+ + H+
MA2+("-*)
(1) (2)
I n the above scheme, higher complexes are favored a t high pH. The values for formation constants determined under many environmental cond i t i o n ~indicate ~ a factor 5 10 for K1IK-2
where K1 and K z are the formation constants for reactions 1 and 2, respectively. I n the present study the rate of first complex formation between Cu2+ and acetylacetone was followed conductometrically. Experimental Materials .-Reagent grade methanol was reflLxed for one hour over magnesium methosidelo and then distilled through a 90-cm. column. The first portion of distillate was discarded, and a second distillation was made over 2,4,6trinitrobenzoic acid through a 15;plate Oldershaw column; b.p., 64.5-65.0', 1it.llvalue 64.75 . C.P. acetylacetone was fractionally crystallized, distilled and the middle fraction collected; b.p. 45" (30 mm.) (unc0r.l; n 1 8 . 5 1.4520 ~ f 0.0001, lit. value 1.45178. Analytical reagent grade Cu-
(1) D. W. Margerum.
(8) J. Bjerrum, "Metal Ammine Formation in Aqueous Solution," P. Ilaase and Son, Copenhagen, 1941. (9) R. Taft and E. Cook, J. Am. Chem. Soc.. 81, 413(1959). (10) A. Weissbrrger, "Techniques of Organic Chemistry," Vol. VII, "Organic Soivants," Interscience Publishers Inc., New York, N. Y.,
(1950).
1955, p. 336.
(7) E. Cook, Thesis. "Rate and Equilibrium Studies with Aqueous TTA." Pennsylvania State University, 1953.
(1933).
(11) G. E. Coates and J. E. Coates. J. A m . Chem. Sor., 06, 2733
