Kinetics of cerium(IV)-perchlorate oxidation of propionic acid

Notes EXPERIMENTAL

Both sequestrants were prepared from 2,5-bis(aminomethyl)tetrahydrofurane or 2,6-bis(aminomethyl)tetrahydropyrane by carboxymethylation with chloroacetic acid. The diamines were prepared by converting respective hydroxy methyl derivative via tosylate into the phthalimide by hydrazinolysis[1,2]. Commercially unavailable 2,5-bis(hydroxymethyl)tetrahydrofurane and 2,6-bis(hydroxymethyl)tetrahydropyranewere prepared by the known procedures [3, 4]. 2,5 - bis(aminomethyl)tetrahydrofurane - N,N,N'N' - tetraacetic acid (THFC) cotourless crystals from water, m.p. 142-144°C, Analysis corresponds to C~M2~OgN2.xH.,O. IR spectrum: 3500(vsI, 3100(w), 1680(vs), 1425(m), 1350(m), 1025(m), 900(s) and 690(s) cm '. 2,6 - bis(aminomethyl)tetrahydropyrane - N,N,N'N' - tetraacetic acid (THPC/ colourless crystals from water, m.p. 230-232°C. Analysis corresponds to C,,H~,O~N2. IR spectrum: 3600(s), 3050(s), 2230(m), 1720(w), 1650(m), 1450(m), 1420(m), 1330(m), 1260(w), 1080(w), 1050(m), 1000(w), 960(s), 920(m), 910(m~, 8801m), 845(s), 790(m) and 705(s)cm '. Lead(II) complexes with the ligands were prepared by the method of Brintzinger[5] and had a general composition Pb(PbL). xH20. According to the elemental analysis, n was 1,3, 3,1 for L = EDTA. BAETA, THFC and THPC, respectively. Stoichiometric stability constants of lead complexes with THFC and THPC as well as the acid dissociation constants of both ligands were determined by potentiometric pH titration at 25°C in 0.1 tool dm ~ NaNO,. The initial estimates of the equilibrium constants (obtained graphically) were refined by using a simplified SCOGS procedure[6l. RESULTS AND DISCUSSION

The pK values of both ligands are given in Table 1. Owing to the considerable spatial separation of the two CH~N(CH2COOH)~ moieties in both THFC and THPC, pK, and pK~ values are close to one another indicating a very weak mutual interaction, if any. The same is observed with pK~ and pK,. In the case of THPC a peculiar relation is found: pK~ > pkg. The possibility of a computational artifact may safely be

ruled out because the slope of the protonation curve was adequately high. During the refinement of the estimates of pK values (guessed from the protonation curve) a reasonable good fit to the original titration curve was obtained. When taking other (arbitrarily chosen) estimates of the pK's the procedure failed to converge if a reverse relation (pK~
Table I. Acid dissociation constants of THFC and THPC and the stability constants of their lead(II) complexes at 25°C. I(NaNO~) - 0.1 mole dm ~ THFC pK~ pK~ pK~ pK~ Ig(K~,,/dm~ mole ') Ig(K,,,,,,jdm mole )

2.42+0.07 2.73+0.05 8.67 + 0.03 8.95+0.01 14.30+0.09 --

THPC 2.71 +0.04 2.98+0.115 9.21 + 0.02 7.96 0.11 14.87+0.04 18.66 + 0.04

~Faculty of Technology, University of Zagreb, 41000 Zagreb, Croatia. Yugoslavia.

! in,~rg ~u~l Chem. 1977, Vo] 39, pp 2095-2097. Pergamon Press.

2095

N. PAULI('? N. IVI(~!I(~ K. JAKOPCI(~t V1. SIMEON O . A . WEBER

REFERENCES 1. F. A. Newth and L. F. Wiggins, J. Chem. Soc. 155 (1948) 2. A. C. Cope and B. C. Anderson, J. Am. Chem. Soc. 77. '995 (19551. 3. A. C. Cope and A. Fournier Jr., J. Am. Chem. Soc. 79, 3',396 (19571. 4. V. G. M. Jones and L. F. Wiggins, J. Chem. Soc. 2 (19451. 5. H. Brintzinger, H. Thiele und U. Muller, Z. Anorg Allg. Chem. 251,285 (19431. 6. I. G. Sayce, Talanta 15, 1397 (19681. 7. L. G. Sill4n and A. E. Martell, Stability Constants of MetalIon Complexes. Spec. Publ. No. 17, p. 673. The Chemical Society, London (1%41. 8. D. T. Sawyer and J. E. Tackett, Z Am. Chem. Soc. 85. 2390 (19631.

Prinled in Great Britain

Kinetics of cerium(IV)-perchlorate oxidation of propionic acid (Received 6 October 1976; received for publication 4 March 19771 The kinetics of the oxidation of various organic compounds by Ce(IV) have been reported by a number of authors. The studies have revealed several mechanisms[l-7] are possible with different reductants, thus conjecture based on analogy would be

naive. The cationic species of ceric in perchloric acid medium have been shown[8-14] to include Ce4÷, Ce(OH) 3+, Ce-O-Ce 6+ and Ce(OH)f +. More recently, Wiberg and Ford[15] have studied the equilibrium among the Ce(IV) species in aqueous

21)96

Notes

acetic acid and have reported that the principal equilibrium is between monomeric and trimeric species. On the basis of our experimental findings we propose that the reactive species is Ce(H20), 4+. Spectrophotometric studies of the Ce(IV)-perchlorate in perchloric acid medium suggests that the predominant species of the Ce(1V) is [Ce(H20),, IOH] 3+ because the absorption intensity in the wave length region 211-+Sm,u. is a maximum compared to the other absorptions at 237 and 252 m#. At higher concentrations of Ce(lV)-perchlorate, the intensity in the wave length region 211-+5 m,a falls off markedly indicating dimerisation. It is gratifying to note that a first order rate expression fits the disappearance of Ce(IV) in perchloric acid medium but the magnitude of k, depends on the initial concentration of Ce(IV)perchlorate (Table 1). The rate of oxidation increases with [H'I and a plot of I/k, against I[[H +] is linear (Fig. 1) and the Michaelis-Menten reciprocal plot of llk~ against l/[propionic acid] is also a straight line with an intercept on the l/kt-axis (Fig. 1). Hence. complex formation between Ce(IV) and propionic acid occurs initially. This has been further confirmed spectrophotometrically by taking the absorption spectra of the reaction mixture in the UV region. The absorption intensity in the range 214-221 mtz is increased, compared to the absorption spectra of Ce(IV)-perchlorate under identical conditions. Addition of NaCIO4 and Ba(C104)2 leads to a positive salt effect. The salt effect is more pronounced in the case of NaCIO4 compared to Ba(CIO4)2 (Table 2). The effect of temperature on the rate of reaction and activation parameters are given in Table 3. From the experimental results and the catalytic behaviour of H + the mechanism can be visualised as below:

Table 2. Effect of addition of salts on the rate of reaction. Temp. 50°C, [ceric perchlorate]=4.00xl0-3N, [propionic acid]= 2.00 x 10 ~M and [perchloric acid] = 1.50 N [NaCIO4] x N k, x 104 sec ~ [Ba(CIO4)2]× l0 M kj × 104 sec 0.00 0.20 0.50 0.80 1.00 1.20 1.50

1.013 1.055 1.138 1.261 1.369 1.458 1.550

0.00 1.00 2.00 3.00 5.00

Table 3. Effect of variation of temperature on the reaction velocity and activation parameters for overall reaction. [Ceric perchtorate] = 4.00 x 10-3 N, [propionic acid] = 2.00 x 10-t M and [perchloric acid] = 1.50 N Temp. °C 45 50 55 60 k I × 104 sec ' 0.533 1 . 0 1 3 1 . 6 7 9 2.935 Energy of activation (AE)=23.72kcal/mole ~, Entropy of activation (AS ~) = -2.38 e.u. and free energy of activation (AF") = 24.48 kcal/mole L

Kz

2[Ce(H~O)._,OH] 3. ~

Ce-O-Ce 6+ + 15H20

24.0

36.0

[Ce(H20). IC2H£OOH] 4+ + H20

[Ce(H,O).] 4+

(1)

t °°l

(5) Although, we have considered three species of Ce(IV), the other species described earlier may also be present. Equation (4) is the rate determining step in which the free radical is formed through C-H bond fission which subsequently undergoes further fragmentation[4] to give rise to products.

//

<,.o; / / ;

where k~ is the observed velocity constant and [Ce(IV)]r is the total concentration of Ce(IV) in the reaction. [Ce(IV)lr = [Ce(H,O), ]4* + [Ce(H=O),,_ IOH] 3+ + Ce-O-Ce ~÷ + K3[Ce(H20),]4*[C2HsCOOH]. o.

~

(6)

-d/dt[Ce(lV)] - kl[Ce(IV)]r

,.°l..°Xti (o,o~

K4

Free radical + [Ce(HzO).] 4+ ----+ product + H + + [Ce(H20)6] 3+.

/ /

°.°I//r/ 04

(3)

[Ce(H20). ,C2H~COOHp+ ~ - ~ free radical + H + + [Ce(H20)6] 3+ ~l,,w (4)

//

80

(2)

K3

[Ce(HzO).I 4+ + C2HsCOOH ~

KI

[Ce(H20) . iOH] 3+ + H ~~

1.013 1.028 1.024 1.038 t.086

,;

I [Propionic ocid]

,.

,.,

2o

(7)

At constant hydrogen ion concentration Ce-O-Ce 6+ formed will be proportional to the [Ce(H20), iOH] 3+ and hence,

A imM -i

Thus, k4K3[H+] [Ce(IV)]r[C2H£OOH] kt[Ce(IV)]r -

+K t

1 + K,[H +] + K3K,[H +] [C2H~COOH] (9)

Table 1. Effect of variation of Ce(IV) on the reaction velocity. Temp. 50°C, [propionic acid] = 2.00 x 10 ~M and [perchloric acid] = 1.50 N [Ceric perchlorate] X 103N k t × 104 sec -I

(8)

Ce_O_Ce~÷ = KI[Ce(H20), IOH] 3..

Fig. l. Plot A(©) shows the effect of variation of propionic acid concentration on the rate of reaction; [Ce(CIO4),d = 4.00 x 10-3, [HCIO,d = 1.50N. Plot B(@) shows perchloric acid effect on the rate of reaction; [Ce(CIO4)4] = 4.00 × 10-3 N, [propionic acid] = 2.00 x 10 ~M.

6.67

5.00

4.00

2.86

2.22

1.67

1.25

1.00

0.959

0.984

1.013 1.053 1.088 1.119 1.167 1.331

2(197

Notes Table 4. Temp. 30°C [Propionic acid] x 103 M

ICeric perchlorate] X 10'}~ N

[Ce(IV)]/[CH3CH2COOH]

1.00

1.00

6.00

1.00

1.25

6.19

1.00

1.67

6.28 Average value of equi. = 6.16

I+K'

Kl ]

1

l/k~=kk4m3tra~]+k4K3-,7~7 [C2HsCOOH]

+K 1

k4"

(10)

According to this, at constant [H+], a plot of l/k~ against I/[C2HsCOOH] will be a straight line, as observed. The plot of I/k~ against 1/[H+] at constant propionic acid will also be linear as experimentally obtained. At lower initial concentration of Ce(IV) the dimer species will be of negligible importance and hence K' in eqn (9) will not appear. Consequently, the value of k~ at lower concentrations will be greater.

Stoichiometry and products A known excess of Cel[V)-perchlorate and fixed amount of propionic acid were allowed to react completely at 30°C and after three months the unreacted Ce(IV) was estimated volumetrically. The stoichiometric ratios are given in Table 4. The product formed was found to be acetic acid which was confirmed by ascending paper chromatography with authentic samples using n-butanol saturated with ammonia as a developing solvent and bromophenol blue as a spot revealing reagent. The Ri value found was 0.60. On the basis of equi,~alents and the product formed, the reaction can be shown to proceed as follows: O C H~--CH~-- C --OH + 6Ce(1V) + 2H20 ---~ CH~--COOH + 6Ce(Ill) + 6H* + CO2

Ill)

EXPERIMENTALANDMATERIALSEMPLOYED Propionic acid used was A.R.(B.D.H.) grade and eerie perchlorate was prepared[16] from the analar quality of B.D.H. eerie ammonium nitrate. Other chemicals used were perchloric acid (60%, Riedel), sodium perchlorate (Riedel), barium perchlorate (E. Merck), sodium hydroxide A.R.(B.D.H.), ferrous ammonium sulphate A.R.(B.D.H.) and ferroin (E. Merck, Grade). l o the known volume of perchloric acid and eerie perchlorate which were kept in a thermostat was added the required volume of propionic acid which was placed in the same thermostat. At suitable intervals, 5.00 ml of this reaction mixture were analysed for cerium(IV) after quenching the reaction with a known

) im~re nu{/ ('item.

amount of ferrous ammonium sulphate and titrating the excess of ferrous ion with eerie sulphate solution, using ferrion as indicator. Reaction bottles were painted black so as to eliminate any photochemical reduction. ']['he rate constants were obtained by plotting a graph of the time against log {a - x); the slope gives the value of k t.

Acknowledgement--One of the authors (I.C.T.) is grateful to the C.S.I.R., New Delhi for financial assistance. Department of Chemistry University of Allahabad Allahabad India

I.C. TFWAR1 S.R. TRIPATH1

REFERENCES !. M. Ardon, J. Chem. Soc. 1811 (1957). 2. J. S. Littler and W. A. Waters, J. Chem. So(. 2767 (1%01: ibid. 832 (1962). 3. G. G. Guelbault and W. H. McGurdy. Jr., J. Phy. (7Item. 67. 283 (1%3). 4. D. Paquetti and M. Zador, Can. J. Chem. 46, 3507 (1968). 5. T. R. Balasubramanian and N. Venkatsubramanian, Ind. J Chem. 9, 36 (197l). 6. B. Krishna and K. C. Tewari, J. Chem. Soc. 3097 (1%1). 7. R. N. Mehrotra, Ind. J. Chem. 10, 1077 (1972}. 8. M. S. Sherilly, C. B. King and R. C. Spooner, J. Am. Chem. Soc. 65, 170 (1943). 9. F. R. Duke and R. F. Parachen, J. Am. Chem. Soc. 78, 150 (1956). 10. M. G. Evans and N. Uri, Nature 166, 602 i1950). 11. F. B. Baker, T. W. Newton and M. Khan, J. Phys. CJ~em. 64. 109 (1960). 12. L. Heidt and M. Smith, J. Am. Chem. Soc. 70, 2476 (1958). 13. T. J. Hardwick and E. Robertson, Can. J. Chem. 29, 818 (1951). 14. E. L. King and M. L. Pandow, J. Am. Chem. Soc. 74, 1%6 (1952). 15. K. B. Wiberg and P. C. Ford, lnorg. Chem. 7, 369 (1%8). 16. A. I. Vogel, A Text Book of Quantitative Inorganic Analysis, p. 316. Longmam London (1%9).

ItJ77.",'el 3% ep. 2097-2098 PergamonPress Printed in Great Brilain

Far-IR spectra of the tetrahaiocobaltate complexes of the morpholinium, piperidinium, 1-methyl- and 2-methyl-piperazinium cations (Received 12 April 1977~ In a framework of a systematic investigation on the influence of the ligands possessing different hydrogen bonding abilities on the coordination geometry of the tetrahalometallates [1,2], we have studied the tetrahalocobaltate complexes of the morpholinium and piperidinium cations and of the l-methyl- and 2-methyl-

piperazinium dications[3]. Their magnetic moments and electronic spectra suggest the presence of complexes having distorted tetrahedral symmetries and their IR spectra (400(~-400cm-~) indicate that the hydrogen bonding interactions are partly responsible for their distortion.