1356
NOTES
They differ from those obtained during the decomposition stage CD of triethylaluminum in the presence of TiCL This is due partly to the catalytic properties of the latter system. The elucidation of the exact mechanism of the decomposition would need further study.’ The analysis of the reaction products points to a carbonium ion mechanism. Were the reaction to pass through free ethyl radicals these would be trapped by the large excess of CC1, solvent and only C2HsC1 would be obtained. Furthermore a free radical decomposition would not explain the specificity of C2HsA1C12;just the opposite, it would be expected that the far more reactive triethylaluminum would react faster than ethylaluminum dichloride. The decomposition therefore must be ionic. Were the decomposition anionic, one would expect mainly the Wurtz reaction product, butane, from Me-C2H6 and the CZHSC~ which is present. Only 1% of butane was detected. Also the carbanionic decomposition product should yield C&&CC&rather than C2HsC1. Tentatively it may be suggested that an equilibrium passing through intermediate steps is set up 2CC14
+ CzHsAlC12 +2CCL.
-t CzHs+AIClr-
The equilibrium is displaced toward the right by recombination of the CClg radicals giving CzCla; hexachloroethane could indeed be detected. It is reasonable to suppose that the ethylcarbonium will react further with both CC4 and CzHbAIClz (a) C2H,+AlC14-
(b) CzHs+AlCla-
+ CC14 +C2&C1 + CCl,+AlC&-
+ CzHsAlClz
CCld ----f
C2Ha
.1
+ AlCla
+ AlCl*-+CzH4AlClg
J.
C2& 4- 2A1Cla
All the major reaction products stripulatedactually have been found. However, the carbonium ion mechanism suggested in this case is in contrast with the carbanionic and free radical mechanisms established for the more normal decompositions of the metal-carbon bond. (7) However, fu:i ther work in this direction is not being undertaken.
OXIDATION OF PHOSPHOROUS ACID BY B. SILVERAND Z. Luz Weizmann Institute of Science, RehoooEh, Iorael Received DeCE?nbsT 8, 1881
The oxidation of phosphorous acid by halogens has been the subject of previous studies. Mitchell’ examined the kinetics of oxidation by iodine in dilute HC1 and concluded that two mechanisms were operative, viz.,the attack of Is-ion on an “active” tautomer of phosphorous acid, and simultaneously a direct attack of molecular iodine on the “normal” form of the acid. He found that the over-all rate decreases with increasing acidity in the range 0.05 to 0.15 N HC1. Berthoud and BergerZstudied (1) A. D. Mitchell, J . Chem. SOC.,123, 2241 (1923). (2) A. Berthoud and W. E. Berger, J . chim. phys., 26, 568 (1928).
Vol. 66
the oxidation of HaPOa by iodine in both acid and neutralsolutions. On the basis of experiments carried out between 0.1 and 0.5 N HC1 they conclude that the oxidation rate is acid independent and involves a direct attack of IZand Is- on phosphorous acid and its ions. In acetate buffer between -pH 4.5 and 5.0 they found a decrease in rate with increasing acidity and, although they fail to analyze their experimental data, put forward a mechanism consisting of the attack of molecular iodine on the mono- or dianion of phosphorous acid, catalyzed by hydroxide ion. Griffith and McKeown3 studied the oxidation up to 1.8 M acid using bromine and chlorine, and for the case of bromine were able to interpret their results as indicating the oxidation of the phosphite ions by molecular bromine, but unlike Mitchell they did not postulate the existence of tautomerism. The above works were reviewed by Yost and R u s ~ e l l ,whose ~ conclusions are that tautomerism plays no part in the oxidation, and that all the results probably are explicable in terms of different reactivities of the two phosphite ions. Van Wazers also comments on the obscurity of the mechanism of oxidation and suggests a chain reaction as one possible mechanism. Stranks and Wilkind state that the absence of measurable deuterium exchange of the phosphorus-bonded hydrogen of H3P03 rules out a tautomeric equilibrium prior to oxidation. In 1959 Martin’ showed that a t high acidities such an exchange does in fact occur, but he did not discuss the relevance of this result to the oxidation reaction. The acid and base catalyzed oxidation of the dialkyl esters of phosphorous acid has been shown to follow the same rate law as that for exchange with deuterium of the phosphorusbonded hydr0gen.8.~ On the basis of this result it was suggested that tautomerism precedes both exchange and oxidation. Considering Martin’s work and by analogy to the dialkylphosphonates it seems likely that tautomerism should play an important role in the oxidation of phosphorous acid in strong acid solutions. In the present work the oxidation and exchange reactions were re-investigated in the strong acid range (up to 4.0 N ) . The reaction also was studied at pH 8.6, where the dominant species of the phosphorous acid is the dianion HPOS--. In the acidic range the results suggest that tautomerism plays a role in the oxidation reaction. At acidities below -1 N tautomerism takes no significant part in the oxidation reaction. Experimental Materials.-Normal and deuterated phos horous arid were made from phosphorus trichloride and $ 2 0 or D?O, respectively.10 The products were purified by crystallization (3) R. 0. Griffith and A. McKeown, Trans. Faraday Soc.. 29, 611 (1933). (4) D.M. Yost and H. Russell, “Systematic Inorganic Chemistry,” Prentice-Hall, Inc., New York, N. Y., 1946. ( 5 ) J. R. Van Wszer, “Phosphorus and i t s Compounds,” Interscience Publishers, Ino., New York, N. Y.,1958, Chap. 7. (6) D. R . Stranks and R. D. Wilkins, Chem. Revs., 67,743 (1957). (7) R. B. Martin, J . A m . Chem. SOC.,81, 1574 (1959). (8) Z. Luz and B. Silver, ibid., 83, 4518 (1961). (9) B. Silver and Z. Luz, ibid., 84, 1091 (1962). (10) I). Voigt and F. Gallais, “Inorganic Syntheses,’’ ed., J. C. Bailar, Vol. IV, McGraw-Hill Book Co., New York, N. Y., 1953.
1357
NOTES
July, 1962 froni HzOor DzOand subsequently dried under high vacuum. NaHaPOa and MazDPQ8 were prepared by otentiometric titration of the corresponding acid with N d H . All other materials were Analar reagent grade. Kinetics. I. Exchange.-The exchange of the phosphorus-bonded hydrogen of HaPOa with solvent DzQ could be observed only In strong acid and was studied by an n.m.r. technique previously described .a The decrease in intensity of the resonance line of the phosphorus-bonded hydrogen was measured using as an internal reference the methyl resonance of methanol which was present as 1% by volume. 11. Oxidation was followed by titration of iodine with standard sodium thiosulfate solution. At p H 8.6, where the oxidation is fast, aliquots of the reaction solution were quenched in sufficient dilute HCl to effectively stop the oxidation.
-*
r-
I
I
I
I
I
+',
Results and Discussion Strong Acid Range.-Martin's7 observation of the acid-catalyzed exchange of the phosphorus4 8 12 bonded hydrogen in HsPO3 was re-investigated at Time hours. lower concentrations of H3PO3. The results are Fig. 1.-The decrease in iodine concentration as a function surnmarized in Table I, from which it may be seen time for the oxidation of phosphorous acid by iodine in that the exchange rate increases with DC1 concen- of 1.0 N HCI; temp. 22". (Compositions of solutions are given tration. I n view of these results and by analogy in Table 11.) with the case of the dialky'phosphonates, the possibility arises that tautomerism plays a role in the oxidation of H3P03 in acidic solutions. If tautomerism is the rate-determining step in oxidation, the oxidation reaction should show acid catalysis and, a t sufficiently high iodine concentrations, should be independent of iodine concentration and have a comparable rate to the exchange reaction a t the same acidity. TABLEI PSEUDO FLRST-ORDER RATECONSTANT FOR EXCHANGE OF THE PHOEIPHORUS-BONDED HYDRQGEN OF A SOLUTION OF 2.4 M HaPOs IN DzO AS A FUNCTION OF [DCll Temp. 22' DCL, mole ] . - I
k X 10. hr.-1
1.49 2.24 3.25 3.60 4.10
1.18 1.49 2.03 2.31 2.82
Concenlratlon
of
ladme.
Fig. 2.-Initial slopes of the plots of decrease in iodine concentration for the oxidation of H3P03 in various acid solutions; temp. 22". (Compositions of solutions are given in Table 11.)
The oxidation of H3P03 by iodine was studied a t €IC1 concentrations of 1.0, 2.5, and 4.0 N HCl, the iodine concentration being varied a t TABLEI1 fixed acidity. The decrease in iodine concentraINITIAL SLOPES OF THE PLOTS OF DECREASE IN IODINE CONtion as a function of time was determined by CENTRATION FOR THE OXIDATION OF 0.05 M PHOSPHOROUS titration with sodium thiosulfate solution. The ACID results for 1.0 N HC1 are shown in Fig. 1. Due The ratio [KI]/[IP] was 4.67 for all runs, temp. 2 2 O . to the fact that the decrease in titer is small comHC1 concn., Initial iodine eoncn., Initial slopes pared to the total iodine titer, the accuracy of the M M mole/hr. X 108 runs is low and no attempt was made to make an 1.0 0.075 0.90 exact kinetic interpretation. However, a number .120 1.13 of qualitative observations can be made by plot,180 1.38 ting the initial slopes of the runs a t different acidi.240 1.62 ties and iodine concentrations against iodine con2.5 0.030 0.75 centration. The results are summarized in Table .075 1.58 I1 and shown in Fig. 2. Two conclusions may be .120 2.25 drawn from the results. First that the oxidation .165 3.13 is acid-catalyzed and secondly that a t 1.0 N HC1 4.0 0.015 0.95 the initial rate is not linearly proportional to iodine .030 1.63 concentration. As may be ween from the last .060 run in Table 11 an appreciable kinetic isotope 3.50 .090 effect results from the substitution of deuterium 4.75 for the phosphorus-bonded hydrogen of HaP03. 1.05 0.240 0.65 Martin observed a similar isotope effect in the case This run is for 0.05 M D-OP(OH)%.
NOTES
1358
of the exchange reaction of H-PO(OH)2and D-PO-
(OH),. The prediction that the oxidation is acid-catalyzed is confirmed by the present results. Previous statements that the rate decreases with increasing acidity1 or is independent of acidity2 are explained by the limited range of acidity studied. The decrease of oxidation rate observed by Mitchell is explicable in terms of a corresponding decrease in the proportion of dissociated phosphorous acid to undissociated acid, since, as mill be discussed later, the ions are oxidized much faster. The expectation that the rate should approach a limiting value a t high iodine concentrations is fulfilled by the results in 1.0 N HC1. For the reasons stated under Experimental it was not possible to study oxidation a t higher ratios of iodine to phosphorous acid, a t which we believe there would be complete independence of rate on iodine concentration. Above 1.0 N HC1 the region of iodine-independent oxidation is presumably a t still higher iodine concentration. The following mechanism is suggested to account for the acidcatalyzed exchange and oxidation. The scheme H.PO(OH)z + H + I_ H*P(OH)3+J _ Oxidation r--------+
O.P(OR)$
:P( 0H)s -
(1)
D.PO(OH)z
L +
Exchange
shows a reversible protonation of the phosphoryl group of H3P03 followed by fission of the phosphorus-hydrogen bond to give a tricovaleiit form of phosphorous acid. This form is expected to be extremely reactive, undergoing oxidation in the presence of m oxidizing agent or accepting a proton (or deuteron) from the solvent to re-form the species H.PO(OH)2. I n the presence of sufficient iodine the exchange reaction cannot compete with oxidation and the oxidation rate becomes independent of the concentration of oxidizing agent. Such an approach to independence a t high iodine concentration may be seen from the results in 1.0 N HC1. From the initial slope of the run at the highest iodine concentration, a pseudo-first-order rate constant of approximately 0.3 X 10-1 hr.-l may be estimated for this run. This value is close to the asymptotic value for 1.0 N HC1, and is in fact of the same order as the pseudo-firstorder rate constant for exchange at the same acidity, i.e., 1.0 x lO-lhr.-l. Oxidation of Phosphite Anions.-Phosphorous acid has two ionizable hydrogens with pK’s of 1.29 and 6.74.5 I n order to isolate the oxidation reaction of the dianion the reaction was studied at pH 8.6 (borate buffer), where the dianion is the predominating species. The experimental solutions contain a mixture of molecular iodine and triiodide ion. It was observed that KI depresses the reaction rate, which suggests that the main oxidizing agent is molecular iodine, since the addition of KI under the experimental conditions decreases the concentration of molecular iodine appreciably but hardly has any effect on the 13- concentration. In fact the
Vol, 66
kinetics could be interpreted in terms of a single significant reaction, vix., the bimolecular attack of molecular I2 on the dianion
ddtm
= -h2[Iz][HP08--]
(1)
Equat,ion 1may be integrated to give S
+ 3b
(a--r;)
In
(F)] (2)
where a and b are the initial concentrations of iodine and phosphorous acid, respectively, 6 is the difference in initial concentration of potassium iodide and iodine, and x is the decrease in concentration of iodine up to time t, as determined by the thiosulfate titration. K is the equilibrium constant for the reaction I 2 I- $ 13-. It should be noted that the derivation of eq. 2 from eq. 1 rests on the assumption that, under the experimental conditions, the concentration of molecular Iz is small compared to both 6 and the concentration of 13-. This assumption follows from the large valuell of K, 7.68 X 102mole-l. Equation 2 implies that a plot of R against time should be linear, with a slope that is independent of 6. Such behavior is in fact followed over several half-lives, as may be seen from the experimental data plotted in Fig. 3. From the slope a value of 7.1 X lo3 mole-l min.-l was derived for k ~ .The value of k2 was unaffected, within experimental error, by doubling and tripling the buffer concentration. The experimental results are consistent with a mechanism involving the attack of molecular iodine on phosphite ions. The possibility that hypoiodous acid (HIO) is the effective oxidizer is ruled out by an experiment carried out in NaOH. Under these conditions the hydrolysis of iodine to form HI0 is practically complete and no measurable oxidation of H3P03was in fact observed. In the case of the dialkylphosphonates a correlation was observed between the rate of oxidation in acetate buffer and the acetate ion-catalyzed exchange of the phosphorus-bonded h y d r ~ g e n . ~ Such a correlation cannot be made in the case of phosphorous acid since we have observed that acetate buffer has no catalytic effect on oxidation and neither is there any measurable exchange of phosphorus-bonded hydrogen with deuterium in solutions of H3P03 in acetate buffer a t pH 5.3 at room temperature. The absence of base-catalyzed hydrogen exchange ’was confirmed by experiments in strong NaOH solution. Martin’ and Brodskii12 have noted the absence of exchange in solutions of Na2HP03 and NaH2P03. The negative charge on the anions of phosphorous acid thus appears to be sufficient to retard attack of a negatively charged base, such as occurs in the case of dialkyl phosphonates. A further feature of the behavior of Hap03 is the fact that an appreciable kinetic isotope effect was observed in the oxidation of the deuterated
+
(11) I. Y. Kolthoff and R . Beloher, “Volumetric Analysis,” Vol. 111, lntersoience Publishers, lno., New York, N. Y., 1957, p. 202. (12) A. 1. Brodskii and L. V. Sulima, Doklady A k a d . .%‘auk S.S.S. R., 85, 1277 (1952).
July, 1962
NOTES
1359
crease in hydrostatic pressure should generally inhibit the formation of micelles and so raise the critical micelle concentration (c.m.c.). To test this prediction, some direct measurements have been made of the influence of pressure on the c.m.c. of solutions of sodium dodecyl sulfate.
T i m e minutes.
Fig. 3.-Oxidation of 0.01 M phosphorous acid in 0.1 borate buffer a t pH 8.6. Initial concentration of iodine 0.015 M ; temp, 22". R (as defined in eq. 2) is plotted vs. time, for different values of 6. The filled dots are for deuterated phosphorous acid, D.PO(OH)z, under the same experimenial conditions.
Experimental The c.m.c. was determined by measuring the specific conductivity of the solutions as a function of the molality of sodium dodecyl sulfate. At each pressure, the points lay on two straight lines whose intersection was taken Go be the c.m.c.7 (see Fig. I). The conductivity measurements were carried out in a Teflon cell8,Qfitted with platinized platinum electrodes and mounted in a conventional steel pressure vessel.1° The temperature of the cell was controlled to within 3~0.01~ and the pressure to within l t l 0 atm. The conductances of the solutions were measured by a Wayne-Kerr B221 Universal Transformer Bridge and were converted to specific conductivities by subtracting the measured conductance of water at the appropriate pressures and by applying cell constants corrected for the linear contraction of Teflon under pressure.9,11 Two different samples of purified sodium dodecyl sulfate gave identical results.
Results Conductivity measurements were made at eIeven concentrations between 2 X and 4 X lo-, acid D.PO(OH),. In Fig. 3 the results of runs molal, at 25") and at pressures of 1, 500, 1000, 1500, carried out with deuterated phosphorous acid and 2000 atni. are plotted together with the data for the normal Figure 1shows the results of some of the measureform. The kinetic isotope effect kH/kn N 3.6. ments. For the sake of clarity in the diagram, the Some runs were carried out in acetate buffer conductivity data at 500 and 1500 atm. have been at pH 4.4 and 5.3, the runs at pH 4.4 being slower omitted. The inset in Fig. 1 shows the variation of than those at pH 5.3 and both sets being appre- the c.m.c. with pressure. The measured value at 1 ciably slower than those at pH 8.6. The decrease atm. (0.00827 mole/kg.) agrees well with values remay be attributed to the decrease in concentration ported earlier by Goddard and Benson' (0.0084 of the dianion, and the much lower oxidation rate niole/kg.) and by Flockhart and Ubbelohdel2 of the monoanion. I n principle it should be pos- (0.0080 mole/kg.). sible to derive a value for the bimolecular constant Discussion for the oxidation of the monoanion. However, The results show that the c.m.c. initially insuch a derivation depends on an accurate knorvledge creases with increasing pressure in accordance with of the dissociation constants of H3P03at the variety of ionic strengths obtaining under the experi- the general prediction made in the introduction. On the basis of Stainsby and Alexander's13 quasimental conditions. Such data are not, available thermodynamic treatment of micelle formation, but a rough estimate gave an upper limit of -10 mole-l min.-l, and it is believed that the reaction and its later refinement by Phillips,14 we should of the monoanion is intermediate between that of expect the influence of pressure on the c.m.c. to be given by the relationship the undissociated acid and the dianion. Investigations were supported in part by a re(a In (c.m.c.)) T,m = A B (1) search grant (RC 5842) from the Division of Reap 1.8 RT search Grants, U.S. Public Health Service. B. S. where R is the gas constant, T is the3bsolute temis the recipient of the nilax and Rebecca Schrire perature, P is the pressure, and AV denotes the Medical Research Grant. change of partial molar volume when the salt passes
THE INFLUENCE OF PRESSURE ON THE FORMATIOK OF MICELLES IN AQUEOUS SO1,fJTIONS OF SODIUM DODECYL SULFATE BY
8. D.
HAMhNN
Divssion of Physical Chemistry Australean Commonwealth Seientzfic and industrzal Research Organ&zntzon, Fishermen's Bend, Melbourne, Australia Receaved December l a s 1Q6i
Density measurements1-6 have shown that the partial molar volumes of long-chain aliphatic sa'fts, dissolved in mater, are greater in the micellar state than in the free ionic state. It follows that an in-
(1) D. G. Davies and C. R. Bury, J . Chem. SOC..2263 (1930). (2) C . R. Eury and G. 8.Parry, ibid., 626 (1935). (3) R. G. Paquette, E. C. Lingafelter, a n d H. V. Tartar, J . Bm. Cham. Sac., 65, 686 (1043). (4) R. J. Tetter, J. Phys. Chem., 51, 262 (1947). (5) K. Hess, W. Philippoff, and H. Kiessig, Kolloid Z., 88, 40 (1939). (6) K. A. Wright and H. V. Tartar, -7. A m . Chem. Soc., 61, 544 (1939). (7) E. D. Goddard and G. C . Benson, Can. J . Chem., 36,986 (1951). (8) J. C. Jamieson, J . Chem. Phys., 21, 1385 (1953). (9) S. D. Hamann and W. Straws, Trans. Faraday Soc., 61, 1684 (1955). (10) J, Buchanan a n d S.D. Ramann. ibid., 49, 1425 (1963). (11) C. E. Weir, J . Reseaich Natl. Bur. Standards, 53, 245 (1954). (12) B. D. Flockhart and A. R. Ubbelohde. J . Colloid Sci., 8 , 428 (1953). (13) G. Stainsby and A. E. Alexander, Trans. Faraday Soe., 46, 587 (1 950).
(14) J. N. Phillips, ibid., 51, 561 (1955).
