NOTES Although we believe our calculations on these and other molecules have demonstrated some interesting general features, the present uncertainties in applying RRKM theory to photodissociation are considerable and have been discussed elsewhere.2
Acknowledgment. This work was supported by the Simmons College Fund for Research. (9) R. J. Campbell, E. W. Schlag, and B. W. Ristow, J. Amer. Chem. Soc., 89, 5098 (1967).
The Oxidation of Hypophosphorous
Acid by Chromium(V1) by J. N. Cooper
955 were removed from samples before analysis by passage through a chloride-form anion column with a large excess capacity. I n the subsequent P(1) determination, the column elu&ntwas heated to 90" for 1 hr in 3 M HC1 purged with Nz to hydrolyze the Cr(1II) hypophosphite. Analysis was performed as usual on the cooled sample. In the P(II1) determination, the sample was allowed to remain in the phosphate bufferiodine mixture for 4 hr to assure complete reaction. The dissociation quotient of H3P02 at 25" was estimated in 1 M LiClOa using a glass electrode and saturated NaC1-calomel reference electrode calibrated against solutions of known hydrogen ion concentration in 1 M perchlorate. The compositions of the solutions studied are summarized in Table I. Over a sevenfold
- ~ Table I: pH of HaPOzSolutions in 1 M LiClO4 at 25'
Department of Chemistry, Bucknell University, Lewisburg, Pennsylvania 17857 (Received August $6,1969)
The report by Haight and coworkers' of a preliminary study of the oxidation of hypophosphorous acid by Cr(V1) in 1 M aqueous perchlorate media prompts us to present the results of our work. In 1 M perchlorate we find, in substantial agreement with Haight, et al., a rate law largely first order in Cr(V1) and hypophosphorous acid, but our results differ in some details. To analyze our data more fully, we have determined the dissociation quotient of hypophosphorous acid and the formation quotient of the Cr(V1)-P(1) complex. The interpretation of data in an earlier study by Pan and Lin2 of this reaction in sulfate media, proposing a rate law second order in Cr(VI), is in error. Apparently the formation of a Cr(V1)-P(1) complex was overlooked; the absorbance data given are clearly first order if their initial absorbance datum is disregarded. Experimental Section J. T. Baker 50% H3P02was purified by the fractional melting t e ~ h n i q u e stored ,~ under nitrogen, and determined iod~metrically.~Unused solutions of HaPOz were discarded 3 weeks after preparation or when they were found to have contained more than 1% H3P03. Recrystallized potassium dichromate was used as a source of Cr(VI), and water was redistilled from alkaline permanganate. Otherwise reagent grade chemicals were used without further purification. All solutions were purged with deoxygenated nitrogen before use. On a 0.01 to 0.5 F scale all of the Cr(II1) product was cationic and was separated from neutral and anionic products using analytical grade exchange resins. Analysis for product chromium was performed at 372 nm following alkaline oxidation with hydrogen peroxide, e(Cr02-) = 4.82 X lo3. Coordinated P(1) and P(II1) were determined by modifications of the standard iodometric procedure^.^ Free H,P02 and H3P03
a
HaPOz, F
Observed p [H+I
6?~(HaPoz),'A4
0.184 0.123 0.0925 0.0615 0.0308 0.0246
1.024 zk 0.004 1.140 f 0.005 1.234 i 0.003 1.362 i 0.006 1.620 i 0.005 1.681 f 0.005
0.099 0,103 0.098 0,105 0.085 0.114
&~(HaPoz) = [Hf][%POz-I / [IlsPOz].
change in formal Hap02 concentration, &~(H3P02)= 0.101 f 0.015. The formation quotient of the Cr(V1)-P(1) complex was estimated photometrically in 1 M perchlorate media, [H+] = 0.161 M ; reduction of Cr(V1) was negligible during the time of measurement. Haight, et d . , l observed no evidence for complex formation between HzPO2- and Cr(VI), and we interpret the data in Table I1 in terms of a 1: 1 complex between the prinTable 11: Absorbance of Solutions of Cr(V1) and HsPOz at 350 nm
[ H f ] = 0.161 A4; temp, 2 j 0 , path length = 2.00
HaPO2, F Asso
x
=
1.00 cm; Cr(S'I)
10-4~
0.000 0.311
0.01016 0.296
0.01695 0.287
0.0339 0.270
0.0508
0.255
cipal species, H3P02 and HCrOd-, corresponding to the anhydride, H2PCr05-. Hypophosphorous acid was in large excess over Cr(VI), and its molar concentration (1) G. P. Haight, Jr., M . Rose, and J. Preer, J . Amer. Chem. Soc., 90, 4809 (1968). (2) K. Pan and S. H. Lin, J. Chin. Chem. SOC.(Taipei), (11) 7, 75 (1960). (3) W. A. Jenkins and R. T. Jones, J . Amer. Chem. Soc., 74, 1353 (1952).
(4) R. T. Jones and E. H. Swift, Anal. Chem., 25, 1272 (1953). Volume 743Number 4
February 19, 1970
NOTES
956 was calculated using the Q~(Hap02)obtained above. Absorbances were measured a t 350 nm where HCr04-, the complex, H2Cr04, and Cr20j2-absorb. The latter two species were minor contributors and the observed absorbances were adjusted using the Qc(HzCr04) = 4.2 for the dissociation quotient of chromic acid and &~(HCr04-)= 98 for the dimerization quotient of bichromate. 3Iolar absorptivities, c(HZCr04) = 925 and t(Cr20j2-) = 3000, were estimated from the absorbance of K2Cr207solutions in 1 M perchlorate. I n all cases the adjustment was less than 5% of the observed absorbance, and the results did not depend critically on the specific molar absorptivities chosen. The formation quotient and molar absorptivity obtained were Qf = 13 =!= 2, e(comp1ex) = 650 f 100. Kinetics were followed at 350 nm on a Beclcrnan DU spectrophotometer equipped with a thermostated cell holder; infinite time absorbances were less than 1% of the initial absorbances. Individual runs, flooding with hypophosphorous and perchloric acids, mere cleanly first order in the absorbance and the pseudo-first-order rate constant was essentially constant over the range F. of initial Cr(V1) concentrations, 1.50 to 8.00 X
Results and Discussion
Table 111: Pseudo-First-Order Rate Constants (25") W+I, M
1.00 1.00 1 .oo 1 .oo 1 .oo 1 .oo 0.607 0.371 0.206 0.131 0.0818 0.632 0.742 0.0569 0.300 0.272 0.260 0.247 0.368 0.325 a
[HsPOz] X IO2, M
4.51 2.71 4.57 9.10 1.83 0.918 4.29 3.94 3.35 2.82 2.24 1.93 3.07 2.62 18.07 8.78 5.21 1.72 47.4 27.6
Median deviation: 1.9%.
k
x
Obsd
106, 8 0 0 - 1
Calcda
21.gb 14.4 22.0 35.0 10.3 5.43 14.8 10.5 7.28 5.43 4,lob 3.44b 5.51b 4.54b 26.2 15.7 11.3 4.45 56.0 36.3
21.6 14.5 21.8 34.9 10.4 5.58 14.5 10.3 7.20 5.49 4.14 3.50 5.34 4.50 27.5 16.4 11 .o 4.33 .55.0
37.1
Average of two runs.
served that for initial H3P02 concentrations greater than about 0.1 M , the rate constants increased with time elapsed since purification of the acid. We used HSP02 within three weeks of purification; Haight, et al., used reagent chemicals without further purification. The mechanism proposed by Haight, et al., involved formation of the anhydride complex, protonation of the complex, and abstraction of a phosphinic proton from the complex by basic species in solution with reduction to Cr(1V). This mechanism is consistent with the 1s present rate with these modifications. The ICO term may correspond either to reaction of the unprotonated rate = [ H C ~ O A - ] [ H ~ P O ~ ]kl[H+] {I~~ complex involving an indeterminant number of solvent Icz[H+I2 J~3[H3P02]1 molecules or to abstraction from the neutral, protonated complex of a phosphinic proton by hydroxide ion. The where second-order rate constant for this latter process would rate = -d[HCr04-]/dt be near to or less than the diffusion-controlled limit if the acid dissociation quotient of the protonated form of IC, = 1 . 8 7 x 10-3, kl = 4 . 9 6 x 10-3, the Cr(V1)-P(1) complex is less than about 10. Direct k2 = 1.61 X = 9.53 X abstraction of a phosphinic proton from Hap02 by HCr04- is also consistent with the kinetics but is [HCrO?-] = [Cr(VI)I/( 1 [H+I/&c thought unlikely because of the relatively weak basicQr [&PO21 4- ~ & D [ H C ~ O1? - ] ity of bichromate ions and the unfavorability of transferring another negative charge to the HCr04- moiety. and The kz term is too large to attribute solely to the changes in activity coefficients typically found in pro[ H ~ P O Z=] [P(I)I/( 1 &A/[H+I] ceeding from 1 F LiClO4 to 1 F HC104,8,9although there The observed and calculated rate constants a t 25" are presented in Table 111. Our rate law was of the same ( 5 ) J. Y. Tong, Inorg. Chem., 3, 1804 (1964). general form as that suggested by Haight, et aZ.,I (6) J. H. Espenson and D. F. Dustin, ibid.,8 , 1760 (1969). but the rate constants we observed were lower except for (7) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Butthose at low [H3P02]. In preliminary work we obterworth and Co. Ltd., London, 1959, p 491.
Our value of & A ( H ~ P O=~ )0.101 in LiC104 js compared with that reported recently for LiN03 media,e &A = 0.135; agreement is good if the mean activity coefficient for a 1: 1 electrolyte in 1 M LiC104 or LiN03 is taken approximately that for 1 M Lie104 or LiN03, re~pectively.~Our value of &t(complex) = 13 is in reasonable agreement with the previous kinetic estimate,'&* = 11. The rate law obtained from 25 observed pseudofirst-order rate constants, k = -d In (A360)ldt (sec-l),
+
+
+
+
+
+
The Journal of Physical Chemistry
NOTES
957
may be some contribution from this effect. Mechanistically, a second-order term in hydrogen ion concentration corresponds to the formation of a doubly protonated complex H4PCr06 +, followed by proton abstraction by water. We have no independent evidence for the formation of such a species, but studies on the exchange of the phosphinic protons in hypophosphorous acidlo show the reaction to be acid catalyzed, suggesting a transient species, H4P02+, whose structure would be analogous to H&rPOs+ formed prior to or during the rate-determining reduction of Cr(V1) in the ic2 step. The principal phosphorus product is P(II1); the ratio, total P(III)/Cr, is 1.55 f 0.05. The Cr(II1) product is green and is not displaced from an acidic cation exchange column by barium perchlorate. The product band is somewhat spread out by cerium perchlorate but Cra+(aq),when added to the column after a sample of the product, passes through the product band under cerium displacement. We interpret this as indicating complex Cr(II1) products with ionic charge greater than or equal to +3.*!" I n all cases both P(1) and P(II1) were found in the cationic product. The reaction is exothermic and the ratios of bound P(1) and P(II1) were somewhat scattered as is shown in Table IV, diminishing with increasing reaction time and the
Product Cr(II1) Complex
0.02
Reaction time, hr
3
0.02
10
0.02 0.2
45 1
Temp, OC
30 30 30 70
+ Cr(V1) -+ 2Cr(V) Cr(V) + H3P02 +Cr(II1) + P(II1) Cr(1V)
Acknowledgments. The author expresses thanks to Professor G. P. Haight, Jr., for his thoughful comments and to the Research Corporation and Bucknell University for grants in partial support of this work. (8) J. H. Espenson and D. E. Binau, Inorg. Chem., 5, 1365 (1966). (9) T. TV. Newton and F. B. Baker, J . Phys. Chem., 67, 1425 (1963). (10) TV. A. Jenkins and D. M. Yost, J . Inorg. Xucl. Chem., 11, 297 (1959); A. Fratiello and E. W. Anderson, J . Amer. Chem. SOC.,85, 519 (1963). (11) J. E. Finholt, K. G. Caulton, and J. W. Libbey, Inorg. Chem., 3 , 1801 (1964). (12) E. Deutsch and H. Taube, ibid., 7, 1532 (1968).
P(IIIj/Cr
1.2 i0.2
1.3 f 0.1
...
1 . 0 f 0.1 0.7 i 0 . 1
of Benzene in Micellar and Norimicellar
0.7 & 0 . 1
Aqueous Solutions of Organic Saltsla
0 . 4 i 0.1
(11)
nonoxidative ligand capture implies that coordination of Cr(V), and perhaps Cr(IV), is as fast as or faster than reduction of Cr(V). On the basis of crystal field energies, Cr(1V) and Cr(V) ought to be substitution labile. Their high formal charges may well encourage rapid and extensive coordination, and steps I and I1 may involve several substituted Cr species.
P(Ij/Cr
..,
(1)
(13) D. K. Wakefield and W. P. Schaap, ibid., 8,512 (1969). (14) This was pointed out by Professor G. P. Haight, Jr. (15) M. T.Beck and I. Bardi, Acta Chim. Acad. Sci., Hung., 29, 3 (1961).
Table IV: Variation of P(1) and P(II1) in the
Reactant concn, F
I n terms of the fast steps, I and 11, presumed to follow the rate-determining reduction of Cr(V1) t o Cr(1V)
elevated temperatures produced at higher reactant concentrations. The reported rate of hydrolysis of Cr(H2P02) 2 + in acidic media8 suggests that the diminution of these ratios with increasing reaction time and temperature may be attributed to partial hydrolysis of the product. The apparent ionic charge of + 3 or greater may be due to the formation of polynuclear complexes; or alternatively, protonated mononuclear complexes, such as have been reported for Cr(II1)acetate12 and Cr(III)-~yanide,'~ may be formed. The weak basicity of hypophosphite and phosphite ions does not appear to support this latter interpretation and the study of the Cr(HzP02)2+hydrolysis showed little evidence for such a species under cation column conditions. The nonoxidative capture of a ligand during the reduction of Cr(V1) has been observed in other ~ y s t e m s , specifically '~ in the reduction with hydrazine in the presence of EDTA.15
Medium Effects on Hydrogen-1 Chemical Shift
by John E. Gordon,lb J. Colin Robertson, and Robert L. Thorne Woods Hole Oceanographic Institution, Woods Hole, Massachusetts 08649,and Department of Chemistry, Kent State Uniaersity, Kent, Ohio 44840 (Receioed A.ugust 22, 1969)
The nature of the medium experienced by molecules in micelles is a subject of considerable interest. An elegant application of nuclear magnetic resonance to this problem has been made by Muller and Birlihahn12who prepared systems containing appropriate fluorine tags in surfactant and solubilized nonelectrolyte molecules and used l9F spectra to study the (1) (a) Contribution No. 2372 from the Woods Hole Oceanographic Institution. (b) Department of Chemistry, Kent State Gniversity, Kent, Ohio 44240. (2) N. Muller and R. H. Birlthahn, J . Phys. Chem., 71, 957 (1967);
72, 583 (1968).
Volume 74, Number 4 February 19, 1970
