Indian Journal of Chemistry Vol. 41 A. February 2002. pp. 270-278

Kinetics and mechanism of permanganate oxidation of phenylphosphinic acid in acid perchlorate solution

Kamla Sharma & Raj Narain Mehrotra*

Department of Chemistry. JNV Univers ity. Jodhpur 342 005. India

Received 24 February 2001; revised 30 October 2001

The stoichiometry of the title reaction is expressed by the equation. Mn04 - + 2C6HsHP(O)OH + 4W ~ Mn(lII ) + 2Cr, l-i sPO(OH h + 2H 20. The kinetics and the rapid scan of the reaction mixture suggest the formation of the complexes by C(, HsHP(O)OH(CI . K!) and C6l-isl-iP(O)0- (C2. K2) with Mn04 - ion . The equilibrium constants of the complexes are W dependent. It is suggested that the complexes are formed by hydrogen bonds. P-H .... O-Mn. and the redox invo lve the di s­sociation of H- from the P-H bond causing transfer of two electrons from substrate to Mn04 - through -0- bridge. The relat ive stability of the complex C2 > CI because KI < K2 and. kl > k2 where kl and k2 are the rate constants fo r the decom­pos ition of the complexes CI and C2. The kinetics in perchlorate solution (11 = I mol dm--\ LiC104) under pse do-first order conditi ons ([PPA] » [Mn04 -n indicated that the W catalysed ra te correlates with the expression kobs(K" + IH;n = a + b[W] + c[H+f where kobs is the pseudo-first order rate constant and Ka is the dissociation constant of C6HsHP(O)OH. The [H+f term is likely due to the acid catalysis of Cl. The activation parameters corresponding to kl and k2 are reported .

Although the oxidation of phosphinic acid has re­cei ved considerable attention t. scant attention has been paid to the oxidation of phenylphosphinic acid (C6HsHP(O)OH, PPA). There is evidence for the ex­istence of the "inactive" (C6HsPH(O)OH) and "ac­tive" (C6HsP(OH)2) tautomers2. The pKa value3

:::-; 1 suggests the possibility of participation of the ionic moiety, C6HsPH(O)O- (A), in the redox reaction. However, C6HsPH(O)O- does not exist in Equi librium with the tautomeric4 form (B). Hence. no ambiguity is expected about the nature of the reacting anion.

o OH

I I C6HS-P-O- + C6HS-P-O-

I H

(A) (B)

Earlier the oxidation of PPA by V(V)s, Cr(VI)6 and FeL/+, (L = bipyridyl or 1,1O-phenanthroline) had been reported7

. In V(V) oxidation, no intermediate was detected and H atom transfer in the r:lte­determining step resulted in the release of a proton and the products. The H+ catalysis was due to proto­nation of V02+ ion and both V02+ and V(OHh2 + re­ac ted with C6HsP(OHh and C6HsHP(O)OH respec­tively. The formation of an ester C04CrH+ OH(O)PHC6Hs ~ H20 + C6HsHP(O)OCr03-), sup-

ported by the rapid scan of the reaction mixture and the kinetics, was considered in the Cr(VI) oxidation . The ester is protonated in the next equilibrium step. The rate-limiting hydride transfer results in the for­mation of PhP+(O)OH that hydrolyses to the product PhPO(OHh in the fast step. In the outer-sphere oxida­tion by FeL/+, an electron is transfen-ed with the si ­multaneous release of a proton. MnO.\- is inert to ester formation and relatively less prone to protonation compared to CrO/- ion. Its redox potential is higher; hence the mechanism for permanganate oxidations are likely to be very different from those of HCrO/ ion . Oxidation of PPA turns out to be a good example of such difference.

Materials and Methods Freshly prepared potassium permanganate (Sarab­

hai-M, GR) solution was boiled, cooled, filtered and standardised spectrophotometricalll (£m = 2227 dm3

mol- I cm- I). PPA (Fluka, practical) was twice recrys­tallised from hot water and dried (m.p. 82°C). The solution was standardised against a fresh alkali solu­tion standardised previously against potassium hydro­genphthalate (BDH, AnalaR). Perchloric acid (E. Merck, GR) was standardised agains t the same alkali solution. Solutions of lithium and sodium perchlorate (G.F. Smith) were prepared by weighing and used to maintain the ionic strength (Il = I mol dm-J

) . Sodium fluoride (BDH, AnalaR) solution was prepared by

SHARMA el at. : KINETICS OF PERMANGANATE OXIDATION OF PHENYLPHOSPHINIC ACID 271

weighing. Doubly di stilled wate r was used through­out.

Delt/eration ofC(H 5HP(O)OH The deuterated phenylphosphinic ac id,

C6H50P(0 )OH , is not com merciall y available. There­fo re, a concentrated solution of C6H5HP(0 )O H in 0 20 was air evaporated. The separated solid was re­dissolved in minimum 0 20 and the solution evapo­rated to dryness. The process was repeated several times to ensure maximum deuteration of the P-H bond. The solution of the deuterated sample was pre­pared in doubly di stilled water just before its use by direct weighing of the sample.

Rate measurements Kinetics were studied under pseudo-first order con-

ditions ([PPA] » [Mn O~ ]) at constant ionic strength

)..l = 1 mol dm-3 (LiCI04). The progress was monitored

at 525 nm (disappearance of Mn04 - ion) using a Un­ion-Giken RA-401 stopped-flow spectrophotometer interfaced with a computer for the acquisition and analysis of the rate data. Temperature of the reactant solutions and reaction cell was maintained using a Haake D8G refrigerated circulatory water bath . The exponential curves for the disappearance of perman­ganate ion were monophasic and the first order plots were linear for more than three half-lives(R2 > 0.998). The pseudo-first order rate constant, k obs, were repro­

ducible within ± 5-7%. The mean k obs values are re­ported in the Tables.

Initially , the rates were measured in the presence of F and P20 7

4- ions, which form stable complexes with Mn3+ ion9

, in order to check the reactivity of Mn3+ ions towards the substrate. Since the increase in the concentrations of either of the ions had no effect on the k obs values, it was concluded that Mn3+ ions did not participate in the rate-determining step. Addition of Ba++ ions, trap for Mn(Y), to the reaction mixture did not precipitate Ba3(Mn04h indicating that the re­duction of Mn(Y) by PPA was probably faster than its precipitation by the Ba++ ions. Therefore, the final study was carried out in the absence of these ions .

Oxidations of PPA and C6H5PO(OHh Oxidation of C6H5PO(OH h , presumed oxidation

product of PPA, by Mn O~ was briefly studied using

HP8452A spectrophotometer. The measured k obs = (2.3 1 ± 0.15)xI0-4 S- I at 30°C indicated that it did not

interfere in the ox idation of PPA (l04[MnO~] = 6.05,

I03[C6H5PO(OHh] = lA, [H+] = 1 mol dm -3).

]n another ex periment, (1O.J[MnO.J -] = 1.8, 104[PPA] = 7 A and 103[F] = 1.0 mol dm -3 at 30°C), k obs was measured at 266 nm (which appeared after the disappearance of Mn04 - peak at 525 nm) after the

reaction was complete. The measured k obs = 9.3x I 0-5

S- I suggested that the stabilized Mn3+ ion did not inter­fere in the rate measurements of the main reaction.

Test for free radical The solutions of PPA and MnO.J- were degassed

with nitrogen and Iml of acrylonitrile was then added to each of the solutions. The polymeri sation of the added acrylonitrile was not noted in the indi vidual solutions and the reaction mixture. It indicated that free radicals did not intervene in the reaction.

Spect rophotomet ry Repetitive spectra of the reaction mixture having

C6H5PO(OH)2 and Mn04 - ion were recorded every four seconds after initiation of the reaction over a pe­riod of one minute using a HP 8452A spectropho­tometer. These spectra were overlaid on the spectra of Mn04 - ion of the same concentration indicating that any complex, if formed between C6H5PO(OHh and Mn04 - ion, was very weak that escaped detection by a reasonably sensitive spectrophotometer.

Rapid scan The rapid scans of the reaction mixture

(l04[Mn04-] = 5.0, 103[PPA] = 5.0 and [H+] = 0.5

mol dm-3 at 25°C) were recorded at the central wave­length 523 nm on a Union Giken RA-401 stopped­flow spectrophotometer connected to a RA-415 rapid scan attachment. A gate time (the period between two successive scanning) and interval (non-working pe­riod for the multi-channel photo-detector unit) of 1 ms were used. Thus the second and subsequent spectra in Fig. I were recorded every 2 ms after the first (low­est) one was recorded after I ms of the mixing of the reactant solutions. The saturation of the spectrum oc­curs in about II ms in comparison to half-lives ex­tending from 200-1500 ms for the redox reaction (104[Mn04-] = 5.0 and [H+] = 0.5 mol dm['3 at 25EC)

for the initial [PPA] = 0.1 and 0.005 mol dm-3. The increase in [PPA] decreased the saturation period whereas an increase in [H+] had no perceptible effect. C, l-i5PH(0)OH and C6H5HP(0)0- ions form the

complex with Mn O~ ion.

StoichiometlY The spectrophotometric titration of a known

Mn O~ soiution by different [PPA] in [HCI04] = 1

272 INDIAN J CHEM. SEC. A. FEBRUARY 2002

Table 1- Dependence of k Oh' on the initial IC6HsPH(O)OH I at different I WI and temperalllres. 10.11 Mn04 - I = 5.n and ~t = 1.0 mol dm-'

101 C(,HsPH(O)OH I 0.05 0. 1 0.2 0.4 0.6 0.8 1.0 (11101 dm-3

) koh< (5-1)

Temp II-n ( 0c) (mol dm-.1) 10 0. 1 0.204 0.305 D.406 0.485 0.5 19 0.538 0.551

0.3 D.255 0.42 1 0.627 0.827 0.926 0.985 1.02 0.5 0.290 0.492 0.753 1.03 1.1 6 1.25 1.31 0.7 0.3 17 0.544 0.848 1.17 1.35 1.46 1.53 0.85 0.342 0.591 0.928 1.29 1.49 1.62 1.71 1.0 0.36 1 0.624 0.986 1.39 1.61 1.74 1.84

15 0.1 0.264 0.4 18 0.590 0.747 0.8 15 0.860 0.884 0.3 0.308 0.532 0.824 1.15 1.32 1.39 1.48 0.5 0.348 0.608 0.967 1.38 1.6 1 1.75 1.85 0.7 0.382 0.673 1.09 1.58 1.84 2.02 2.14 0.85 0.404 0.7 13 1.1 5 1.67 1.96 2. 16 2.29 1.0 0.43 1 0.762 1.24 1.81 2. 12 2.33 2.48

20 0. 1 0.33 1 0.552 0.838 1. 12 1.27 1.36 1.44 0.3 0.368 0.643 1.03 1.48 1.72 1.88 1.98 0.5 0.4 10 0.725 1.1 8 1.73 2.04 2.24 2.39 0.7 0.446 0.798 1.32 1.95 2.32 2.57 2.74 0.85 0.477 0.854 1.42 2. 12 2.54 2.8 1 3.0 1 1.0 0.510 0.9 14 1.52 2.28 2.74 3.04 3.25

25 0. 1 0.383 0.674 1.06 1.42 1.77 1.91 2.03 0.3 0.421 0.751 1.24 1.83 2. 19 2.4 1 2.58 0.5 0.461 0.833 1.39 2.09 2.5 1 2.8 1 3.0 1 0.7 0.505 0.914 1.54 2.33 2.83 3.15 3.37 0.85 0.538 0.977 1.65 2.51 3.02 3.36 3.65 1.0 0.57 1 1.04 1.76 2.68 3.27 3.66 3.94

30 0. 1 0.453 0.8 10 1.33 1.97 2.35 2.59 2.77 0.3 0.501 0.906 1.54 2.33 2.83 3.1 5 3.39 0.5 0.542 0.990 1.69 2.60 3. 16 3.56 3.84 0.7 0.586 1.07 1.82 2.83 3.47 3.9 1 4.23 0.85 0.624 1.14 1.96 3.04 3.73 4.20 4.55 1.0 0.664 1.22 2.09 3.25 4.00 4.5 1 4.89

mol dm-3 were carried out at 525 nm at which the in­termediate valence states of Mn are transparent. It was presumed that the presence of C6HsPO(OHh, the im­mediate oxidation product of PPA, had no effect on the measured absorbance over a period of 30s because the rate of its oxidation is several orders slower than

that of PPA. The plot of absorbance/([Mn 0; ]/[PPA])

against ([Mn 0 ; ]I[PPA]) ratio, Fig. 2, indicated that

MPPA]/MMn 0 ; ] = 2. Hence, the stoichiometry of

the reaction is given by Eq. (Equation 1).

Mn O~ +2C6HsHP(0)OH+4W ~Mn(IIl)

... ( I )

Oxidation product The oxidation product of PPA is presumably

C6HsPO(OH)2 though its characterization was not at-

tempted. The presumption is believed to be correct because the oxidation product of phosphinic acid, H2PO(OH), is phosphorous acid HPO(OHh in the corresponding oxidation.

Characterisation of reduced Mn( III) state The reaction mixtures, in the absence of F and

P20 74

- ions, at the complete reduction of Mn 0 ; ion

appeared pinkish-red perhaps due to stabilization of trivalent manganese forming complex either with PPA (present in excess) or C6HsPO(OHh, the oxida­tion product. The oxidation state of manganese at this stage was determined by estimating the liberated io­dine at 352 nm (C352,1O = 19400 dm3 mort cm- I

). A

number of reaction mixtures (l04[Mn 0 ; ] = 2 mol

dm-3) were prepared in which [PPA]/[Mn 0 ; ] ratio

SHARMA et (I I.: KI ETICS OF PERMANGA NATE OX IDATION OF PH ENYLPHOS PHI IC AC ID 273

Fig. I- The successive rapid sca ns o f the reac ti on mi xture ( 104rMn04-1= 5.0, IOJrpPA1= 5.0, rWl = 0.1 and f1 = 1.0 mol dm-J at 20°C) recorded at 523 nm using a 3.5 nm slit. A gate time and interval time of I ms were used. The absorbance is on an arbitrary scale. The lowest spectrum was recorded I ms after mi xing and the subsequent spectrum moving upward were re­corded after every 2 ms. The saturation in the absorbance oc­curred in about II ms after mixing.

,....., v

0 0·2

~ --,....., -< ~

0'16 C :>< 0)

l 0·12 Vl

.I:> -<

o 0-5 1'5 2 2·5 3 3'5 4

Fig. 2-The plot o f absorbance x rC6H5HP(O)OH ]I[Mn04-1 against[Mn04 -]l[C6HsHP(O)OH ]) rati o. The absorbance was measured for the remaining Mn04 - ion in the reaction mixture as soon as the reaction is over.

varied from 2 to 4. Addition of a slight excess of solid KI to reaction mixtures , after permanganate colour had disappeared, resulted in the liberation of the io­dine. The estimated iodine, (9.9 ± 0.2)x I0-s mol dm-3

,

was in fair agreement with the expected value 2x I 0-4

mol dm-3 ([Mn(VIl)] = [Mn(III)] = 0 .5[1z]) confirm­ing that Mn(VJI) was reduced to Mn(III).

The pinki sh-red colour of the reaction mixture, af­ter permanganate ion was completely reduced, decol­ourised over a period probably due to the slow reac­tion shown in Eg. (2). The slow nature of the reaction

o 0.02 0.04 0.06 0.08 0.1

Fig. 3-The nonlinear plots o f k obs against [C6H5HP(O)OHl at

[HCI041 = 0.1 mol dm-3 and temperatures 10° (0). 15° (. ), 20°

(L'1), 25° (0) and 30°C (0 ).

(2) was confirmed by the visually observed slow di s­appearance of the pinkish-red colour that developed on mixing Mna/ + and PPA in [HCI04] ~ 4 mol dm-3

.

Hence, the rate measurements were not attempted.

C6HsPH(0)OH + 2Mn3+ + H20 ---7 C6HSPO(OHh + 2Mn2+ + 2H+ (2)

Results Dependence on {C6H 5HP(O)OH]

The results in Table I indicated that the kobs in­creased with increase in [PPA] though the increase is not proportionate. The plots of kobs against [PPA] , Fig . 3, are nonlinear whereas the plots between kobs-

I

and [PPAr l, Fig. 4, are linear with intercepts. Hence,

the empirical rate law is expressed by the Eg. (3) . It was further noted that b depended on [H+].

k _ k[PPA] obs - a + b[PPA]

... (3)

Dependence on {H%] The results for 1.0 ~ [H+] ~ 0.1 mol dm - 3, Table I ,

indicated that the k obs increased with increase in [H+]. The nonlinear plots of kobs against [H+], Fig. 5, are indicative of a different correlation, which is di s­cussed in the Discussion Section.

Discussion The stoichiometry of the reaction has indicated that

two mol of PPA are oxidized by one mol of Mn 0 ;

ion which is reduced to Mn3+ ion. The reduced Mn3+

274 IND IAN J CI-IEM. SEC. A. FEBRUARY 2002

2.8

2.4

2

~

'" 1.6 -;-

D

1.2 0

-'<

0.8

0.4

o 20 40 60 80 100 120 140 160 180 200 220

lC6H~PH(O)OH r' (dm3 mol")

Fig. 4-The linear plots o r k"hs- ' against [Cr, H, HP(O)OH r' with

intercepts at 25°C wi th [H I j = 0.1 (0). 0.3 (_ ), 0.5 (t.), 0.7 (0).

0.85 ( A ) and 1.0 mol dm-J (D).

state is estab li shed by the estimation of the liberated

iodine from the reaction mixtures in which Mn O~ is

completely reduced . Thus Mn O~ ion is reduced by

four electrons coming from two mol of PPA that are oxidized. Therefore, each PPA mol is providing two electrons. Mn,,/+ is a strong oxidising agent than

Mn O~ yet the oxidation of PPA including its oxida­

tion product C(iHsPO(OHh is extremely slow com­

pared to the Mn O~ oxidation, probably because

Mn,,/+ forms very stable complexes with these sub­strates. The formation of an intermediate complex by

Mn O~ with the substrate is indicated both by the

presence of a Michaelis-Menten type of kinetics and the rapid scan of the reaction mixture. The apparent overall equilibrium constant, denoted by b in Eq. (3), is dependent on [H+] indicating the presence of a protonation-deprotonation equi librium prior to the rate-determining step.

The K" value of permanganic acid is 1.78x I02 dm3

mol- I indicating that HMn04 is a strong acid. There­fore, its deprotonation equi librium is unlikely to in­tluence either the dependence of the rate or the equi­librium constant(s) on [H+] . Thus, the deprotonation of PPA is the other alternative to explain the H+ de­pendence of the rate and the equilibrium constants of the complexes likely to be formed by both C6HsPH(O)OH and C6HsPH(O)O- with Mn04 - ion .

The inability of Mn04- to form ester, similar to those

3.5

0

3

2.5

';" 2 II)

on .0 0

1.5 """

o 0.2 0.4 0.6 0.8

[ H' J ( mol dm-3)

Fig. 5-The nonlinear plots o f k ubs against [H+] 011

[C6H, HP(O)OH] = 0.4 mol dm-J and temperatures 10° (0). 15°

(_ ).20° ( A ), 25° (D) and 30°C (0).

formed in chromic acid oxidations, leads to the sug­gestion that the complexes are formed by hydrogen bonding (P-H···O-Mn). The likely structures of the complexes formed by C6HsPH(O)OH and C6HsPH(O)O- are shown in the alternate mechanism and differ by a proton. In view of the above, the prob­able mechanism of the reaction is likely to involve the following reactions.

K" C6HsPH(O)OH ~ C6HsPH(O)O- + W

KI Mn04 -+C6HsPH(O)OH ~

[02MnOr ·HPOH(O)C6Hsr (el)

K2 Mn04 -+C6HsPH(O)O- ~

[02Mn02···· HPO(O)C6Hs]2- (e2)

kl [02MnOr··HPO(O)C6Hsr --+ H2Mn04-

. . . (4)

... (5)

... (6)

+ C6HsP02 • . . (7)

k2 [02MnOr .. HPO(O)C6Hsfr --+ HMnO/r

+ C6HsP02 . . . (8)

SHARMA el (/1. : KINETICS OF PERMANGANATE OXIDATION OF PHE YLPHOS PI-II NIC ACID 275

kit [02MnOr·· HPOH (O)C6Hsr + H+ ~ H2Mn04-

+ C(\HSP02 + H+ . .. (9)

fast C(\ HSP02 + H20 ( ) C6HsPO(OH h ( 10)

fas t Mn(Y) + C(\HsPH(O)OH + H20 ~ Mn(lll)

+ C(\HsPO(OHh + 2 H+ ( II )

Scheme 1

The dissoc iat ion of P- H bond in the rate deter­. . . d k I-Ilk D h k D mining step IS supporte by ' obs obs = 3.4. T e obs

was measured from the deuteriated sample prepared

Table 2- Va lues of the intercepts (I) and slopes (S) o f the pl ots

between kohs-I and [PPAr l and the calculated values o f (Ka +

11-1+]) x l iS = b (e mpiri cal Eq . 3) at different [WJ and te mpera­tures

10

15

20

25

30

[WI (mo l d m-J )

0. 10

0.30

0.50 0.70

0.85

1.00

0. 10

0.30

0.50 0.70

0.85

1.00 0. 10

0.30

0.50

0.70

0.85

1.00

0.10 0.30 0.50 0.70

0.85

1.00

0.10 0.30 0.50

0.70 0.85

1.00

Intercept

(s)

1.422

0.822

0.622 0.522

0.464

0.427

0.786

0.538

0.418

0.352 0.33 1

0.302 0.552

0.387

0.31 3

0.266

0.240

0.221

0.390 0.284 0.236

0.207 0. 191

0.175 0.264

0.204

0.176 0.160

0.147

0.136

Slope b

(mo l dm-J s)

0.0 162 20.4

0.0 155 2 1.2 0.0 14 1 26.5

0.0132 3 1.7

0.0 123 35.9

0.0117 40.1

0.0140 14.2

0.0 135 16.0

0.0123 20.4

0.0113 25 .0 0.0107 29.4

0.0101 32.9 0.0122 9.5

0.0112 13.8

0.0106 17.7

0.00988 21 .5

0.00929 24.5

0.00871 27.9

0.0111 7.0 0.0105 10.8 0.00966 14.6

0.00887 18.6 0.00834 2 1.8

0.00788 24.4

0.00972 5.4

0.00896 9.1 0.00834

0.00774 0.00728

0.00685

12.7

16.6 19.2

21.9

in the laboratory, which had not been tested for 100% conversion . The kl-llkD = 4.3 was observed in the oxi­dat ion of H}P02 by Mn04- ion and kl-llkD = 2.5 ± 0 .5 was obtained for the decomposition of [Ag(O Hk ·H2-

P02] 2- intermed iate, and kHlkD = 8.2 ± 0.2 for the OW assisted decomposition of the intermediate in the oxi ­dation of H2P02- ion by Ag(OH)4- ion l2. Thus, it

would be appropri ate to assume the P-H bond fissi on in the rate-determining step of the present reacti on.

The rate law, based on the reactions (4)-( I I ), is g iven by Eq. ( 12) and is different from those obtained in the oxidations by vanadium(y)4 and HCr04 - oxi­dationss.

k = (k 2K2Ka +k I K1[H +]+k H K 1[H +]2) [PPA] obs

{Ka +[H +]+(K 2 K a + K1[H +])[PPA]}

( 12)

The poss ibility of the additional [H+]2 term arising from the acid catalysis due to C6HsPH(O)OH is inad­mi ss ible because of the absence of a second order term in [PPA]. The other obvious alternative is to consider the acid catalysis of the complex Ct , which has been included in the proposed Scheme. One common pathway involves ac id catalyzed tautomeri­sation, which will expose a pair of electrons on phos­phorus atom making it an effective two-electron re­ducing agent (or oxygen atom acceptor). Such an "ac­tive" phosphorus atom could be present or produced

by (P-H ~ :O-Mn ~ ) proton transfer in each of the terms proposed for the rate law in Eq. (12). Yet, an­other poss ibility is the direct hydride ion transfer.

The empirical rate law in Eq. (3) is identical with the rate law in Eq. (12), with a = Ka + [H +], b = (K2Ka + K1[H+]), and k = (k2K2Ka + kIKI[H%] + kHKI[H%]2).

The term (K2K. + KI[H+]) in the denominator of Eq. (12) is consistent with the dependence of the equilib­rium constant K, on [H+]. The various constants ap­pearing in Eq. (12) are separated as described below.

Separation of the cons/ants Equation (13), the reciprocal form of the Equation

(12), is in agreement with the reciprocal plots in Fig. 4 where k is defined in Eq. (17).

.. . (13)

276 INDIAN J C1-1 EM. SEC. A. FEBR UA RY 2002

40

35

30

~ 25 J$.. I~

is 20 + ~

~ 15

10

5

0

0 0.2 0.4 0.6 0.8

Fig. 6--The linear plo ts o f (K" + [W])x I/S against IH+]. where I and S are the intercept and slope va lues from the plo ts between

knhs-I and IC6 H, HP(O)OH r l. at temperatu res 10° (0). IS° (0).

20° ( . ), 25° (~) and 30°C (0 ).

The Equi valence o f the intercepts (I ) and the slopes

(5 ) o f the linear plots be tween kobs- ' and [PPAr', Fig. 4, is expressed by Eqs ( 14) and (IS) respectively. The re latio ns gi ven by Eqs (16) and ( 17) are deduced from Eqs ( 14) and (15).

Intercept(l ) = (K2Ka + K,[W])/k

Slope (5 ) = (Ka + [H+])/k

(Ka + [W])x(I/S) = K2Ka + K, [W]

(14)

( IS )

( 16)

(Ka + [H+])/(S) = k = k2K2Ka + k,K,[W] + kHK,[W]2

... (17)

The linear plots of (K, + [H+])x(VS) aga inst [H+], Fig. 6, are consistent with Eq . ( 16). The inte rcepts and slopes of the plots were used to compute the values of K2 and K, respec tive ly. The polyno mial Equation of second order in [H+], Eq . ( 17), is solved for the coef­ficients k2K2Ka, k, K, and kHK, values. The estimated va lues o f the Equilibrium constants K2 and K, helped to obta in the va lues o f k2, k, and k,., since Ka values are known.

Eq. ( 17) could be rearranged to Eq . (1 8) and the linear plots of (k- k2K2Ka)[H+r ' against [W] are in Fig . 7. Using a least squares program, the intercept (=k,K,) and s lope (=kHK ,) values of the plots were calcul ated and used to compute the values o f k, and k,., respecti vely at di ffe rent temperatures. These va lues were in close agreement with the values o btained fro m the soluti on of po lynomial Equati on (17).

. 130 ~

c

~ ~ "" ::s

90

70

o 0.2 0.4 0.6 0.8

Fig. 7-The linear plots be tween (k - k2K2Kj )[H+r l aga inst [WI wi th in tercept (= kI K I) and slope (= kHK I) at Ivmperatures 10° (0 ).

IS° (_ ), 20° (0). 25° (~) and 30°C ( 0 ) where k = [kob,x{K" + [W j + (K2K" + KI[H+])[PPA]} HPPAr l.

... (1 8)

The mean calcul ated values of k" k2 and kH and those of K, and K2, from the two esti matio ns, are in T able 3 wherein are also g iven the cOITesponding ac­tivatio n and thermodynamic parameters. It is inter­es ting to note that K2 > K, indi cating complex CI is less stable (weaker) than C2. Thi s is simi lar to the observati on made in the ox idation of phosphinic acid " and led us to think of a s imi lar Hr transfer from PPA to Mn(VlI ). Ho wever, H+ catalys is of redox and weakening o f the complex argues against such a ra­tionale. The fact that O-bridg ing and expanded coor­dinati on number to 5 o r 6 is commo n fo r phosphorus and also cf metal ions in high ox idation state leads to consider the poss ibility that - P-O mi ght coord inate to the positi ve Mn(VIJ) center or coordinat ion of -Mn-O to the positive P center. This is a lso ruled out on the considerati on that the corresponding ox idat ion o f C6HsPO(OHh, capable o f forming such coord ina­ti on, is extremely slo w; perhaps for the reason that it

is devoid of P- H bond. Hence, the formati on of P-H .. ·O- Mn hydrogen bonds, shown in complexes Cl and C2, appears to be the li keliest fo rmulation fo r the complexes. The redox would requi re di ssociatio n o f H+ from the P- H bond and two-e lec tron transfer

th rough the - 0- bridge from PPA to M n(VlI).

SHARM A 1'/ 01.: KIN ETICS OF PERMANGA ATE OXIDATION OF PHE, YLPHOS PHINIC ACID 277

Table 3- Va lucs o f the rate cons tan ts k, and k2 (S- I). k" . k' and e (dm 1 mol 's- '). e (dn/' 11101- 2 S- I) and the cquilibriulll constants K,

and K~ (d m' mol ') at d iffe rc l1ttc l11pe ra tu rcs and thc va lucs of rc latcd acti v:llion and thc rmodynami c paramcte rs.

TCl11p (OC) 10 15 20

Constanl

k, 1.9 I. ± 0.07 2.94 ;!: 0.05 3.00;!: 0 .1 6

k~ 0.45 ± 0.03 0 .71 ;!: 0 .03 1.1 7 ;!: 0.09

kll 1.09 ± 0.06 1.76;!: 0 .05 2 .23;!: 0.09

k' 55.9;!: 2. 1 63.4;!: 1.1 72.4 ± 3. 1

k2 64 ± 5 75 ± 3 89 ± 7

e 3 1.7 ± 1.9 38 .1 ± I 44.3 ± 2. 7

K, 29 .2 ± n.R 24.1 ± 0.5 20.5 ± 0 .4

K2 142 ± 6 86 ± 3 73 ± 3

6 H' . b.H" (kJ 11101- 1)

-r.. U I I b.5 . b.5 (J K- 11101- )

A/temate mechal/isll/

The data cannot di stingui sh whether complexes formed between PPA and Mn(YII) are involved in the rate determining steps or s imply constitute side reac­tions. Thus a di stinction between internal ox idation­reduction in complexes Cl and C2 and simpl e bimol­ecul ar collisions of C6HsPH(O)O- and C6HsPH(O)OH

with Mn04- ion cannot be made as shown below since the two approaches are likely to have the same acti­vated state.

The rate law derived from the consideration of the above reactions is given by Eq. (22) which is identical with Eq. (12) with kl = klKJ, e = k2K2, and k3 = kHK1•

... (22)

It is proposed that the complex Cl decays by pro­ton transfer from C6HsPH(O)OH to Mn04r exposing a lone pair susceptible to two electron transfer to give

H2Mn04r and C6HsPO(OHh and the complex C2 decays by hydride transfer to HMn042r + C6HSPO(O­Hh and . The W catalyzed decay of complex Cl oc­curs by proton plus two-electron transfer. In the case of bimolecular collisions, the redox would occur by direct hydride transfer from PPA to Mn(YII) as in Equations (19) and (20).

25 30 M I 65 '

4.07;!: 0.04 5. 15;!: 0. 11 3 1 ;!: 4 - 114 ± 14

1.92 ;!: 0.04 2.97;!: 0 .11 65 ± 2 -4 ±3

2.6 1 ;!: 0.03 3.0 1 ;!: 0.09 32;!: 5 -1 12 ± I&

79.4 ± 0.7 94.& ± 1.9 16 ± I - 139 ± 4

96 ± 2 107 ± 4 16 ± I - 137±5

50.9 ± 0 .6 55.3 ± 1.7 18 ± I - 137 ± 4

6 ft" 65"

19 .5 ± 0 .2 I R.4 ± 0 . 1 - 17 ± 2 -30 ± &

50 ± I 30 ± I -47±3 - 125 ± I I

Activation parameters The formation of Cl and C2 complexes is exo­

thermic because the respecti ve 6.H 0 values are nega­tive. The relatively more negative value of

6.H ~ compared to 6.H 1° indicated the relatively

higher stability of the complex C2, which is further substantiated by the fact that K2 > K1• Further, the

relatively more positive 6.H:2 than 6.H : is consi s-, tent with the fact that k2 < kl values. The relatively

more negative 6.S2° value is likely because the com­plex C2 is likely to be more hydrated because of larger negative charge on it compared to the Cl com­plex. The transition state of the complex C2 is likel y to be looser than that of the complex Cl and the same

is reflected by the more positive 6.$:2 value com-

pared to 6.S:' value. It is interesting to note that

6.H:', "" 6.H :' and 6.S:2 "" 6.S:' because kH "" kl .

The two are of the same order and differ very little . The enthalpy and entropy values corresponding to bimolecular constants kl, e and e are almost similar indicating that the transition states in each case is about the same irrespective the charges involved. Al­ternatively, it might be due to the same kind of elec­tron transfer, which as proposed earlier is the hydride

transfer. It is expected that Hr approaching Mn04-along a reaction coordinate experiences strong attrac­tion but distorts the coordination sphere canceling effects of enthalpy of activation, making negative en­tropy the dominant energy barrier to the activation process .

278 INDI AN J C HEM. SEC. A. FEB RUA RY 2002

Acknowledgement Authors are thank ful to the Department of Science

and Technology for the fin ancial ass istance toward the purchase of the stopped-fl ow instrument and Jun­ior Research Fell owship to KS . Thanks are al so due to the UGC for the award of Project fe ll owship to Manu Mehrotra who carri ed out several additi onal experi­ments. Thanks are also due to Pro f. G P Haight for certain suggestions.

References I Haight G P. J 1', Rose M & Preer J . J Alii ehelll Soc. 90 ( 1968)

4~09: Sengupta K K. Pal B B & Mu khe rjee D C. Z Phys Chelll . (Nellj'o lge). 72 ( 1970) 230: Gupta K S & Gupta Y K, J chelll Soc (A). ( 1970) 256 ; Cooper J N. Hoyt H L. Buffing­ton C W & Holrnes C A. J phys Chelll. 7 1 ( 1971) 891; Vi ste A . Ho lm D A, Wang P A & Veith G D, Illorg Chelll , 10 ( 197 1) 63 1; Cooper J N. J phy:; Chelll , 74 ( 1974) 659 ; Gupta K S & G upta Y K, I ll org Ch elll , 13 ( 1974) 853: Mohan D. C hhabra V K & Gupt a Y K, J ehelll . Soc. Dalton Trails ,

( 1975) 1737: Nagori R R, Mehta M & Mchro tra Raj N. J ('helll Soc. Daltoll Tra ils. ( 1979) 2 16: 'l nd rayan A K. Mi shra S K & G upta Y K. Illorg Chl'lII . 20 ( 1981) 450 & 192.+: Ziiho nyi-Bud6 E & Si miind i L I. 11I(1I~r; chill i Acta Lett. 18 1 ( 199 1) 149.

2 Fra nk A W. J org Chelll . 26 ( 1961) 850. 3 Van Wazer John R. Phosp/lOnt.l' alld its COIIII)(JII /uis. Vol

( Int e rsc ience Publ ishers. New York) 1958, p 349. 4 Q uin L D & Dysart M R. J org Chelll . 27 ( 1902) 1012. 5 Mehro tra Raj N. Call J Chelll , 63 ( 1985) 663. 6 S harma K & Mehrotra Raj N. Trails III l't Chelll , 14 ( 1989) 48. 7 S harma K. Prakash A & Mehrotra Raj N. 8111/ ehelll Soc

Japall . 62 ( 1989) 4009. 8 Jaky M. S imandi L I. J chelll Soc, Perkill TI'IIII .I', 2 ( 1976)

939. 9 Dav ies G. Coord chelll Rev. 4 ( 1969) 199. 10 Ford-Smith M H & Rawsthorne J H, J chelll Soc (A), ( 1969)

160 . 11 Sharma K, Mehrotra Raj N & Hai ght G P. I lldiall J Chelll.

39A (2000) 709. 12 Mehrotra Raj N & Kirschenbaum LJ. illorg Chelll, 28 ( 1989)

4327.