The kinetics of Mn(II)-catalysed ozonation of oxalic acid in aqueous solution

Wet. Res. Voi. 26. No. 7. pp. 917-921. 1992 Printed in Great Britain. All rights reserved

0043-1354/92 $5.00 + 0.00 Pergamon Press Ltd

THE KINETICS OF Mn(II)-CATALYSED OZONATION OF OXALIC ACID IN AQUEOUS SOLUTION R. ANDREOZZIt, A. INSOLA~, V. CAPRIO 2 and M. G. D'AMORE2 'lstituto di Ricerche sulla Combustione, CNR, P. le V. Tecchio, 80125 Napoli and 2Dipartimento di lngegneria Chimica, Facolt:i di Ingegneria, Universit:i di Napoli. P. le V. Tecchio, 80125 Napoli, Italy (First receired February 1991; accepted in ret,isedform January 1992) Abstract--The Mn(ll)-catalysed ozonation of oxalic acid in aqueous solution is investigated at pH = 0 and pH = 4.7. Experiments arc performed in both semi-batch and batch conditions by following the kinetic development of the ozonation processes. At pH = 0 first-order kinetics with respect to both ozone and Mn(ll) are observed with no dependence upon the oxalic acid concentration. A rate constant of 6.2 × 1041tool-I rain-I and equimolar consumptions of ozone and oxalic acid are evaluated. At pH =4.7 higher initial reactivities are observed whereas the low ozone to oxalic acid molar consumption ratios indicate the intervention of molecular oxygen in the oxidation process thus confirming its radical evolution. Reaction mechanisms are proposed to explain the observed reactive behaviours of the investigated systems.

Key words--oxalic acid, ozonation, kinetics, manganese catalyst, mechanism, water treatment

INTRODUCTION

Ozonation treatments of wastewaters for organic pollutants abatement are performed by either direct ozone application (Bailey, 1978) or combined action of ozone and systemscapable of enhancing its reactivity (Glaze and Kang, 1989). Advanced ozonation processes, as the last processes are called, make general use of u.v. and/or ozone decomposers to promote radicalic reaction mechanisms thus allowing the destruction of even more stable organic pollutants. In this respect investigations could reveal the existence of a new class of advanced ozonation processes in which the catalytic effect of metal ions, such as Mn(ll), which undergo easy oxidation by ozone to valency states effective towards organic oxidation could be exploited. Only very scant and disconnected information is available from literature concerning this (Tyupalo et al., 1974; Tyupalo, 1981a, b) and no quantitative evaluation of reaction kinetics have since been reported. This investigation is performed in order to achieve the kinetic characterization of the Mn(II)-catalysed ozonation of oxalic acid in aqueous solution. It represents a step towards the more general objective of exploring the opportunities offered by the combined Mn(II)/Oj system for the oxidation of less reactive organic substrates in aqueous solution. EXPEEIMENTAL

Oxidation experiments at 20°C are performed in batch and s~..m/-batchconditions. For semi-batch experiments a 0.3 dm3 glass vessel is used as the ozonation reactor. 0.2 dm3of the aqueous solution of WIL 2&7--D

917

oxalic acid (1.4 raM) and MnSO4(!.2 x 10-3-3.3 x 10"s M) are added to the reactor. All the runs at pH ,-0 are performed in i M HCIO~ aqueous solutions. For runs at pH-.4.7 phosphate buffered solutions are used. An ozonized oxygen stream with ozone content of 3% per volume is flowed through the solutions at a flow rate of 36 dm3/h. Ozone is dispersed into solution by means of a sintered glass disc whereas mechanical stirring ensures the complete mixing of the reacting system. The reaction progress is followed by quantitative evaluation of oxalic acid at different ozonation times. Analysis is performed by means of a high performance liquid chromatograph (HP 1090L), equipped with a C 8 Spherisorb reverse phase column, usin8 a buffered aqueous solution (4 ml H3PO4 85%, 50 ml methyl alcohol in 1000ml aqueous solution). The concentration of the gaseous stream leaving the reactor is continuously monitored during the ozonation process by means of u.v. measurements at 253 nm. Batch ozonation experiments are performed by using a 3 emJ u.v. quartz cell (optical length- l cm) as the reaction chamber. Measured amounts of aqueous solutions containing known concentrations of oxalic acid and ozone respectively are rapidly mixed in the cuvette and submitted to absorbance measurements at 260nm for ozone determination. Oxalic acid evaluation by u.v. measurements at 210nm has also been attempted with poor results because of the ovcrlappin8 of optical signals likely due to the manganese--oxalatecomplexes of a different nature. RESULTS AND DISCUSSION

Results of semi-batch experiments are shown in Fig. I. No appreciable reactivity is observed for direct attack of ozone on oxalic acid whereas the addition of small amounts of Mn(ll) brings the system reactivRy to remarkable levels. After an initial transitory stage, linear decays of oxalic acid up to its complete disappearance are observed during each ozonation

918

R.A.,,,~P.EozzI

et

al. 20

2°t

I .~q

o

~

1.0

x

o

o o

0.5

0~

i

0

l

10

i

i

20

I

I

30

I

I

0

40

I

~

20

40 Time

T i m e (rain)

Fig. I. Ozonation of oxalic acid at pH = 0. I , Oxalic acid, no Mn(ll) addition; I-1, oxalic acid, [Mn(ll) h = 3.3 x 10-5 M; @, oxalic acid, [Mn(ll)lo = 1.2 x 10.3 M; A, ozone in the outlet gaseous stream, [Mn(ll)h = 3.3 x 10-3 M; and A. ozone in the outlet gaseous stream, [Mn(ll)]0= 1.2 x 10-~ M. experiment with slope values increasing at increasing Mn(ll) concentration. For the same experiments the ozone concentrations in the outlet gaseous stream are plotted in Fig. I as a function of the ozonation time. Diagrams indicate that during oxalic acid oxidation, after rapid attainment of stationary conditions, the outlet ozone concentrations remain at constant values thus suggesting a corresponding constancy of the ozone concentrations in the liquid bulk. Ultraviolet measurements, performed on samples of reacting mixture withdrawn from the reactor at different ozonation times, confirm the presence of ozone in the bulk of the reacting solution. The high system reactivity, however, prevents the quantitative evaluation of the actual ozone concentrations of the reacting system because of the time required for the sampling procedure. The presence of dissolved ozone should exclude the possibility that constancy of ozone absorption rate is due to only diffusional limitations of ozone transfer from gas to liquid phase with no implication of reaction kinetics. The stationary ozone concentrations of Fig. 1 appear to be affected by only Mn(lI) contents of the reacting solution with no dependence upon oxalic acid concentration. Results of semi-batch ozonation experiments thus indicate that no influence upon the kinetics of catalytic oxidation is exerted by oxalic acid and that linear trends of Fig. I are the consequence of pseudo-zero-order dependence upon ozone concentration. Diagrams of Fig. 1 also indicate that during each ozonation experiment the molar ozone consumptions are the same as the corresponding molar oxafic acid consumptions. The influence o f p H change upon oxidation process is investigated by means of the experiments of Fig. 2.

I

I

60

80

100

(rain}

Fig. 2. Ozonation of oxalic acid at pH = 4.7, [Mn(ll)~ =

3.9 x 10-s M. O, Oxalic acid, [03] (in the inlet gaseous stream) = 9 . 0 x 10-4 M and @, oxalic acid, [O~](in the inlet gaseous stream) = 1.5 x 10 -'~ M. The increase of pH from 0 to 4.7 results in very different development of the oxidation process. Higher oxidation rates are observed in the initial ozonation stage, a constancy of the outlet ozone concentration is still recorded but linear decay of the oxalic acid concentration is lost since its consumption rate decreases at increasing ozonation time. Contrary to what is observed at pH = 0 the ozoneto-oxalic acid (AO) molar consumption ratios (Fig. 3) do not remain constant during the oxidation process. In the initial ozonation stage the molar consumption ratios are remarkably smaller than unity whereas they rapidly increase with further development of the ozonation process. The decrease of the ozone concentration in the gaseous stream increases the tendency of the oxidation process to keep the ozone-to-oxalic acid consumption ratios at small values. 30

2.0

<~

1.O

o

1

2

3

AA0(mmol)

Fig. 3. Ozone vs oxalic acid molar consumptions at

pH =4.7. O, [03] (in the inlet l~eOUS stream),, 9.0 x 10-4M and @, [03] (in the inlet i~seous stream)= 1.5 × 10-' M.

Ozonation of oxalic acid

919

where k2 ~' k~. This gives the following rate equation for oxalic acid oxidation:

-80

- I/2 d[Mn(llI)-AO]/dt -- - d[AO]/dt = k,[Mn(III)-AO].

(1)

-100

If one assumes that production of the Mn(IIl)/ monoxalate complex arises from slow reaction steps of Mn(II) oxidation by ozone:

u {:

k]

-120

2Mn(II) + O3 + 2H +

,2Mn(III) + HzO + O2

regulated by a first-order kinetic dependence upon Mn(ll) and O , then: -IG

0

i

o

i

I

Ioo

I

2oo

i

I

i

3oo

4oo

- d[O3l/dt = - 1/2 d[Mn(III) - AO]/dt

Time (s) Fig. 4. Ozonation of oxalic acid at pH = 0 batch conditions.

O. Ozone. [AO]o=I0-4M, [Mn(ll)]o=0.1xl0-SM; O. ozone. [AO]o= 10-2M, [Mn(ll)]o=0.1 × 10-SM; A. ozone. [AO]o= 10-4 M. [Mn(ll)]0 = 0.5 x 10-s M; A, ozone. [AO]o= 10-2M, [Mn(ll)]0=0.5 × 10-~M; I1. ozone. [AO]o= i0 -4M. [Mn(II)]u= 1.0 × l0 -sM; I-'1. ozone. [AO]0= l0 -2 M. [Mn(ll)]o = 1.0 x 10-s M and "A', ozone. [AO]0= 10-' M. [Mn(ll)]o = 2.0 x l0 -5 M.

= k, [Mn(ll)] [05 ].

(2)

From equations (I) and (2), after attainment of stationary conditions, the ozone and oxalic acid consumption rates will be: - d[Oj]/dt = - d[AO]/dt = k3[Mn(ll)][O~]. T w o assumptions can be made for Mn(I[) oxidation to Mn(lll) by ozone:

Results of kinetic investigations performed at pH = 0 on a homogeneously reacting system are shown in Fig. 4. The diagrams reveal a first-order dependence of reaction kinetics upon the ozone concentration with no dependence upon the oxalic acid concentration. A linear correlation between slope values of the logarithmic plots of Fig. 4 and the Mn(ll) concentration is also observed. The kinetic development at pH = 0 of the catalytic ozonation of oxalic acid is thus described by the following equation:

Mn(ll) + O 3 +

H +

, Mn(lll)

Mn(ll) + OH" + H +

+ O H " + O2

(slow)

, Mn([ll) + H z O

(fast)

or k3

Mn([l) + O5 + 2H +

, Mn(IV)

- d O 3 / d t -- - dAO/dt -- k[Mn(ll)][O3]. Results reveal that oxalic acid ozonation promoted by Mn(ll) ions proceeds according to different reaction mechanisms whose occurrence is regulated by the adopted pH values. Kinetic behaviour exhibited at pH = 0 in both batch and semi-batch ozonation experiments agrees with the assumption that oxalic acid ozonation occurs by means of a slow reaction step of Mn(II) oxidation by ozone followed by fast reaction steps which cause oxalic acid oxidation at the expense of species deriving from Mn(II) oxidation. It is known from Taube's investigations (Taube, 1948) that addition of Mn(IIl) ions to an aqueous solution of oxalic acid, at pHs favouring the Mn(III)/monoxalate complex, determines fast oxidation of the organic acid according to the following mechanism: Mn(IIl) - AO Mn(llI) + AO"

ks

k2

, Mn(II) + AO ~ , Mn(II)+2CO2

+ 02 + H20 Mn(IV) + Mn(lI)

,2Mn(lll).

(slow) (fast)

The first assumption is invalidated by Height'S experiments of Mn(ll) oxidation by ozone (Nowell and Height, 1987) unless Mn(II)/oxalate complexes behave differently from that found by Hoign6. The second assumption, partly supported by literature data (Tyupalo, 1981a, b), is more consistentwith our experimental results since it agrees with the appearance of a transitory stage for oxalic acidcatalysed ozonation (Fig. I). The transitory stage would correspond to the increase of Mn(IV) concentration up to the stationary value:

[Mn(IV)] =

k3/k4[03]

with a corresponding increase of the oxalic acid oxidation rate. Ozonation experiments performed using aqueous solutionsof perchloricacid previously contacted with M n O 2 give furtherconfirmation of the proposed mechanism since they result in the disappearance of the initialtransitorystage (Fig. 5).

R. A.'qDtEOZZJet al.

920

O

18

OH" +

A O =-

+ 2H +

//

12

, CO: +

H:O .c//+ \ OH

O

"C + O3

' CO2 + OH" + 02

\OH N

OH" + O~

o

o06

HO2, O ; + O~ i

I

5

0

|

I

I0

t

I

15

O; + H ÷

I

, HO] + O2

P H ÷+O~" , O2 + O; , OH" + O:

20

Time (rain) Fig. 5. Ozonation of oxalic acid at pH = 0 , [Mn(ll)lo -- 1.2 x 10-3 M. 0 , Oxalic acid and C), oxalic acid (solution previously contacted with MnO2).

//

O

O

"C + O 2

---OH

\OH O

For runs at pH = 4.7 the observed ozone-to-oxalic acid molar consumption ratios (AO3/AAO < 1) indicate the intervention of oxygen in the whole oxidation process thus suggesting that the ozonation of oxalic acid proceeds through a radicalic mechanism. The same suggestion also arises from the increase of the ozonation efficiency at the smaller ozone concentration since competition by dissolved oxygen becomes more marked. It must be remarked that at the adopted pH any initiation of radicalic ozone decomposition can result in the activation by OH radicals of chain propagating steps not interrupted by HO2 radicals formation (Staehelin and Hoign~, 1982). It must also be considered that contrary to runs at pH = 0, in experiments at pH = 4.7 oxidative development through Mn(Ill)/oxalate complexes is less effective. At this pH Mn(III)/dioxalate and trioxalate complexes are mainly formed which have been reported (Taube, 1948) to be much less reactive than the monooxalate complex. This is also confirmed by direct observations in this investigation since no appreciable tendency to oxidation is exhibited from oxalate after its addition to aqueous solutions of Mn(lll) at p H - 4.7. Consistent with the previous statement that no production of OH radicals can directly derive from the Mn(II) ozonation, the activation of radicalic ozonation mechanisms can probably be ascribed to the decomposition of Mn(III)/oxalate complexes. Accordingly, the following reaction scheme can be proposed for the ozonation process at pH -4.7: Mn(II)

, Mn(III)

Mn(III)(AO2-).

, Mn(II) + A O " + (n - I)AO 2-

AO" + O 3 + H +

° 2CO2 + O2 + OH"

-o.

,

CO2 + HO~

CONCLUSIONS

The refractory behaviour of oxalic acid towards ozone is overcome by the addition of Mn(II) ions to the reacting aqueous solutions. Results of ozonation experiments indicate that Mn(ll) ions exert a different catalytic role at, respectively, pH = 0 and pH = 4.7. At pH = 0 the oxidation of Mn(lI) to Mn(IIl) via Mn(IV) can be assumed as the rate controlling step of the overall oxidation process. Accordingly, first-order kinetics with respect to both ozone and Mn(II) ion concentrations and no dependence upon the oxalic acid concentration are observed. The appearance of an initial transitory stage and its disappearance after contacting of the aqueous solution with MnO2 are also in agreement with the proposed mechanism. At pH = 4.7 a radicalic development of the ozonation process is indicated by the remarkable extent of oxalic acid oxidation at the expense of molecular oxygen. Accordingly a radicalic ozonation mechanism is proposed by which Mn(IIl) oxalate complexes act as radical chain initiators because of their capability to oxidize oxalate ions to oxalate ion radicals. Work is in progress to investigate the opportunities offered by Mn(II) catalysis for the ozonation of other organic substrates. REFERENCES

Bailey P.S. (1978) Ozonatlon in Organic Chemistry. AQtdemic Press, New York. Glaze W. H. and Kang J. W. (1989) Evaluation of the ozone-hydrogen peroxide procets in a temi-batch reactor using tetrachloroethylene as a model compound. In Proceed/ngs of the Ninth Ozone World Congress, New York, pp. 596-617.

Ozonation of oxalic acid Noweil L. H. and Hoign~ J. (1987) Interaction of iron (II) and other transition metals with aqueous ozone. In Proceedings of the Eighth Ozone World Congress, Zurich, pp. ES0-E95. Staehelin J. and Hoign~J. (1982) Decomposition of ozone in water, rate of initiation by hydroxide ions and hydrogen peroxide. Env/r. Sci. Technol. 16, 676--681. Taube H. (1948) The interaction of mansanic ion and oxalate. Rates, equilibria and mechanism../. Am. Chem. Soc. 70, 1216-1220.

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Tyupalo N. F. (1981a) Catalytic oxidation of 4-methylpyridine in sulfuric acid. Zh. priM. Khan. $4, Part 2, 338-340. Tyupalo N. F. (1981b) Study of the iron (II) oxidation by ozone. Doki. Akad. Nauk. $SSR 894-896. Tyupalo N. F., Yakobi V. A. and Bernashevskii N. V. (1974) Oxidation of 2-methyl-5-ethylpyridineby ozone in an acid medium. Ukr. khim. Zh. 40, Part 7, 74--77.