ELSEVIER Chemical Engineering and Processing 35 (1996) 443-449

Ch.mi~Ingin..rlng

andProcessing

Oxidation kinetics of Fe(II)NTA to Fe(III)NTA by oxygen

D. van Velzen*, L.-O. Actis-Dato, J. HofeleCommission 0/ the European Communities, Joint Research Centre, Environment Institute, 21020lspra (VA), Italy

Received 31 October 1995; accepted I3 March 1996

Abstract

The oxidation of aqueous solutions of Fe(II)NTA was investigated in a bubble column reactor. The rate of the oxidationreaction was found to be proportional to the partial pressure of oxygen and to have an order of 0.5 with respect to the Fe(II)NTAconcentration. The overall transfer of oxygen appears to be predominantly reaction controlled for Fe(II)NTA concentrations inthe order of 1 mmol l " I, whereas for high concentrations (100 mmol l ' ') mass transfer becomes the controlling factor. Overallmass transfer coefficients and reaction rate constants were determined in the range between 25 and 60°C and for a wide rangeof Fe(II)NTA concentrations (5-250 mmol 1- I).

Keywords: Oxidation; Kinetics; Bubble column reactor

1. Introduction

During the last decades a number of wet scrubbingprocesses for the removal of nitrogen oxides from wastegases has been investigated. Nearly all of these pro­cesses use a metal-chelate complex to enhance the lowsolubility of NO in aqueous media. The most promisingmetal chelate complexes are the ferrous complexes,

EDTA:

HOOC - CH1' /CH1 -COOHN-CH1-CH1-N,

HOOC - CH1/ CH1 -COOH

NTA:

CHZ-COOH

/

/

N _ CHHOOC'\.. I 2

CH z COOH

• Corresponding author.

0255-2701/95/$15.00 © 1996 - Elsevier Science S.A. All rights reserved

PII S0255-2701 (96)04153-0

Fe(II)EDTA (ethylenediamine-tetracetic acid) andFe(II)NTA (nitrilotriacetic acid).

The large majority of the processes consist of simul­taneous NOx/S02 removal based on an absorption-re­duction mechanism [1-4]. Some processes have beendeveloped up to the pilot plant scale, but as yet suchdevelopment has not yielded a valid alternative to theexisting industrially applied denoxing processes.

Another possibility was recently investigated by vanVelzen et al. [5], who proposed a process where theabsorption is carried out in the absence of S02 using anabsorption-desorption system. Nitrogen oxide is recov­ered by steam stripping in a concentrated form, suitableto be used as a valuable raw material (e.g. as a processgas for hydroxylamine production).

It was shown that the equilibrium constants for theformation of the complex Fe(II)EDTA .NO are ap­proximately 25% higher than those of Fe(II)NTA .NO[6,7]. However, given that there is a considerable pricedifference between the two, Fe(II)NTA is an interestingalternative.

Nearly all flue gases contain a certain percentage ofoxygen. During the absorption step, some of the Fe(II)complex is inevitably oxidised to Fe(III). The Fe(III)chelate complexes are considerably less effective at ab­sorbing NO than the ferrous chelate complexes. It isthus necessary to include a reduction step to maintainthe Fe(II) concentration. This can be easily done byelectrolysis [4,5,7].

444 D. van Velzen et al. / Chemical Engineering and Processing 35 (1996) 443-449

A detailed knowledge of the rate of oxidation ofthe Fe(II) complex is necessary for the preliminarydesign and for the calculation of the energy consump­tion of an electrolysis cell for the regeneration of theused Fe complex solutions.

2. Literature review

In 1952 Jones and Long [8] investigated the stabil­ity of Fe(II) and Fe(III)EDTA complexes. They re­ported that the ferrous complex is relatively unstablewith respect to both dissociation and to oxidation tothe ferric complex. They observed that rapid oxida­tion takes place in Fe(II)EDTA solutions with con­centrations of 4, 10 and 12 mmol 1- I which wereexposed to air under a minimum of agitation: 65%,40%, and 35% respectively of the Fe(II) in the solu­tion reacted to Fe(III) in 8 min. The reaction rateswere not further investigated.

In 1978 Hasui et al. [9] investigated the oxidation ofFe(II)EDTA in a wetted wall column and found thatthe oxidation rate was 1st order with respect to thepartial pressure of oxygen and 0.7th order with respectto the Fe(II)EDTA concentration in the liquid. Anempirical relationship was presented including the liq­uid flow rate through the column. Hasui did not at­tempt to develop a more general rate expression.

Neumann and Lynn [10] studied the oxidation ofFe(II)NTA in a small wetted wall column of 6 mminner diameter in relationship to the development of aprocess for the oxidation of H2S. Air was used as theoxidant gas. They report that the presence of NTAcatalyses the reaction of oxygen and Fe(II). A fast,irreversible, diffusion controlled chemical reaction wassaid to take place. They determined overall masstransfer coefficients which ranged from 7.5 x 10- 5 to1.2 x 10- 4 m s - I for gas flow rates of 0.8 to 2.0 ms - I. They considered the system a gas absorption ofoxygen enhanced by chemical reaction and calculatedthe corresponding enhancement factors, which variedfrom 53-63 depending on the concentration of theFe(II)NTA complex. This concentration was variedbetween 20 and 70 mmol 1- I. Also these authors didnot propose rate equations.

Sada et al. [II] investigated the kinetics of the oxi­dation of Fe(II)EDTA and Fe(II)NTA by experi­ments in a bubble column. They varied the oxygenconcentration from 2 to 7 vol.%, the Fe(II)EDTAconcentration from 5 to 20 mmol 1- I and the tem­perature from 20 to 60 °C. The results were evaluatedon the basis of a rate equation including a term forthe physical mass transfer and a term for the reactionkinetics. The volumetric mass transfer coefficient(PLa) was determined separately by absorption ofoxygen in a sodium sulphite solution. They found ex-

perimental values for PLa of 0.036 to 0.046 s - 1 at 50°C for their bubble column of 1000 ml liquid with aG3 gas sparger of 10 mm diameter. The earlier obser­vation of Hasui et al. that the reaction was 1st orderwith respect to the oxygen concentration was confir­med. The kinetics of the oxidation rate was found tobe 0.5th order with respect to the Fe(II)EDTA con­centration. An excess of 20% EDTA reduces the oxi­dation rate by up to 30%. The rate constant wasdescribed by an Arrhenius-type equation:

k=A exp( ;~a)

where A = 1.09 X 104 mol " " (l-m S)-I with m =0.536 and - E; = - 23.3 kJ mol- 1.

The oxidation of Fe(II)NTA was studied at onetemperature (50 0C). Sada et al. reported that theoxidation rate is 1st order with respect to oxygen and0.7th order with respect to the Fe(II)NTA concentra­tion. Contrary to EDTA, excess NTA does not sup­press the oxidation. Sada et al. observed a decreaseof the pH during the oxidation of Fe(II)NTA,whereas with Fe(II)EDTA the pH increases.

Closer study of Sada's data shows that at 50 °C thereaction is predominantly diffusion controlled for com­plex concentrations higher than 5 mmol 1- 1. When theconcentration is smaller than 0.5 mmol l" 1 the reactionkinetics decrease considerably and the overall oxygentransfer rate becomes reaction controlled.

In 1989 Zeise [12] studied the oxidation ofFe(II)EDTA in a stirred reactor with air as oxidisinggas. He monitored the formation of the Fe(III)EDTAphotometrically (wavelength, 400 nm) and polaro­graphically. He found that in 10-12 min more than90% of a 9 mmol 1- I solution was oxidised. Zeise didnot attempt to determine rate constants.

From the foregoing it can be concluded that themost valuable work on kinetics of the oxidation offerrous EDTA and NTA complexes was done bySada et al. [II]. They investigated extensively the sys­tem Fe(II)EDTA and did limited research onFe(II)NTA. The authors also observed that thescheme of the oxidation reaction for Fe(II)NTA isprobably different from that of Fe(II)EDTA chelate.Therefore, a detailed investigation in the oxidation ofFe(II)NTA is as yet outstanding and will be pre­sented in this paper.

3. Experimental

The experiments on the oxidation of aqueousFe(II)NTA solutions were carried out in a laboratoryscale bubble column containing 300 ml of aqueoussolution. A diagram of the experimental set-up is rep­resented in Fig. 1.

D. van Velzen et al./Chemical Engineering and Processing 35 (1996) 443-449 445

by-paso

(I)

(3)

252010 15t [min]

5

t

4

6

2

oo

10

8C Fe(II)[mmol/l]

The process implies the absorption of oxygen fromthe gas into the liquid phase, followed by the oxidation.The reaction rate depends on the volumetric, overallmass transfer coefficient:

rAcA=c1--pLa

where CA and cl are the oxygen concentrations in thebulk of the liquid and at the interface, and ~La is thevolumetric, overall mass transfer coefficient.

The absorption is followed by a chemical reaction inthe liquid phase. In the literature [9-11J there is ageneral agreement that the reaction rate is proportionalto the oxygen concentration. However there is someuncertainty over the reaction order n with respect to theFe(I1) concentration. The following equation holds:

r; = kRcACB = kRCB(C1- rA ) (2)PLa

where subscript A refers to oxygen, the subscript B tothe Fe(II)NTA complex.

The oxidation of the Fe(II)NTA complex has thefollowing stoichiometry:

4Fe(II)NTA2+ + 2H20 + O2

~4Fe(III)NTA3+ +40H-

One mole O2 oxidises four moles Fe(II), therefore:

rB = 4rA

4. Results and discussion

Fig. 2. Example of the decrease of the Fe(lI) concentration with time(PA =6.1%, CNTA = 10 mmot I" ', gas flow =230 I h -I).

or

Thus, it follows that:

1 dt 1 1-=-= +---r B dCB 4c1kRcB 4c1PLa

eooler

bubble colum..

mUDgenair

The bubble column was operated continuously withrespect to the gas phase and batchwise with respect tothe liquid phase. An air/nitrogen mixture of knowncomposition was bubbled continuously through theFe(II)NTA solution. The gas inlet consists of a series of24 holes of 2 mm diameter. The gas flow rate wasvaried between 100 and 250 I h - 1. The oxygen concen­tration covered a range between 2 and 15 vol.%. Typi­cal flue gas concentrations lie between 2 and 9 vol.%.The oxygen content was measured by an industrialinfrared photometer (model URAS IOE) of the firmHartmann and Braun. Before passing through the reac­tor, the gas was led through two bubble columnscontaining water. In this way the dry gases from the gasbottles were saturated with water vapour which avoidedevaporation losses from the reactor solution. The ex­periments took place at temperatures of 25-60 DC.

For each experimental condition (temperature, oxy­gen concentration, gas flow rate and initial Fe(II)NTAconcentration) 2 I of Fe(II)NTA solution containingequimolar parts of Fe(II) and NTA were prepared. TheFe(II)NTA concentration was varied between 10 and250 mmol 1- 1. To prevent reaction with atmosphericoxygen the solutions were prepared and kept undernitrogen atmosphere.

300 ml of the solution was filled into the reactor andgas passed through for a given period of time. In thisway a series of five or six experiments were carried outwith various gas contact times, e.g. 5, 10, 20, 30 and 45min. For each experiment the initial and final Fe(II)concentrations were analysed by titration with standardammonium metavanadate solution using sodiumdiphenylaminosulphonate as a redox indicator. Forcontrol, the total iron concentration in the solution wasalso determined at the beginning and the end of eachexperiment. In this analysis all iron was reduced withhydroxylamine to Fe(I1) and the IR absorption of thephenanthroline Fe(I1) complex was determined with aphotospectrometer (wavelength, 510 nm).

By this operation for every experimental condition, acurve was obtained of Fe(I1) concentrations as a func­tion of time. An example is given in Fig. 2.

Fig. I. Experimental set-up for the determination of the kinetics ofthe oxidation of Fe(Il)NTA to Fe(II1)NTA.

446 D. van Velzen et al.] Chemical Engineering and Processing 35 (1996) 443-449

The oxygen concentration at the interface, cA, is as­sumed to be in equilibrium with the oxygen partialpressure in the gas phase and Henry's law applies:

LErrErr=-J- (6)

j

The data and the corresponding diagram of a plot of(4PA)/(HrO) versus l/cs, where n = 0.5 are given inTable 1 and Fig. 4.

Rearrangement of Eq. (3) using Henry's law to describecA yields:

~=_l_+_l_ ~HrB kRci PLa

where PA stands for the partial pressure of the oxygenin the gas flow and H for the Henry coefficients foroxygen in water (788 bar 1 mol " ') at 25 "C),

It follows that a plot of (4p~/(HrB) against l/(ci)yields a straight line where the slope is l/kRand the axisintercept l/(~La). The values of the unknown constantswere obtained on the basis of Eq. (5). In this equationthe only unknown parameters are the order of thereaction n, k R and ~La.

Six series of experiments were selected, all carried outat the same temperature (25 "C), gas flow rate (210 1h -I), with O2 concentrations between 4.2 and 4.8 vol.%and with a wide range of Fe(II)NTA concentrations(5-250 mmol I-I). For every experiment the averageoxidation rate and the average Fe(II)NTA concentra­tion were determined.

To determine the order of the reaction, n was variedbetween 0.1 and 1 and the relative error was calculatedfor each regression. The relative error (defined in Eq.(6» as a function of n is given in Fig. 3. It follows thatthe minimum relative error is obtained for n = 0.5.

(7)

0.50.40.30.20.1

..----t

~/

.>:/

~ IJkR

tA:PL a

oo

0.2

PA (bar) CB (mmol l" ') rB (mmol (l min-I»

0.042 204.0 0.5760.042 77.4 0.5330.042 43.1 0.4390.042 9.46 0.3290.042 4.30 0.2510.048 4.29 0.273

kR=0.0137[(~01)-0.S S-IJPLa = 0.060 [s- I]

1

0.4

From the slope and intercept it follows that l/kR=1.22 (mmol 1- 1)-0.5 min and 1/(~La)=0.278 minwhich give the following values for the reaction rateconstant and the volumetric, mass transfer coefficient at25°C:

(4PAYlH I'll)

[min] 0.8

Table 1Data used for the determination of k R and PLa

0.6

The oxygen transfer is predominantly reaction con­trolled for Fe(II)NTA concentrations of the order of 1mmol 1- I. For concentrations of 10 mmol 1- I thecontribution of reaction and diffusion are approxi­mately equal, whereas for higher concentrations (100mmol Ir ') the oxidation becomes practically diffusioncontrolled.

The differential Eq. (3) can be solved analytically.The resulting solution is:

H(

t<I - n) cO<I-n) t 0)CD - D CD - CD

t=- +....;;;...-..;.;.4PA (1- n)kR PLa

The found values of kR , ~La and n can be introducedinto the solution of the integral. The decrease of theFe(II) concentration with time can be computed andcompared with the experimental values obtained duringthe series of experiments. In Fig. 5 the experimentalvalues of three series are represented together with the

(4)

10.80.60.40.2

I

4' I4

2

oo

[%]

Err 6

8

10

exponentn

Fig. 3. The relative error as a function of n. Fig. 4. Determination of k R and PLa.

D. van Velzen et al. / Chemical Engineering and Processing 35 (1996) 443-449 447

8040 60Tee]

/....17

V..... V

---t>

L--'

8

4

6

2

o201008040 60

t[minl20

. ...........-..........

" '-- -- -.:;:: r----h-

I---...

----,- I--r--~ -..., ~-. -.-.o

o

60

30

120

e Fe<m 90[mmollll

Fig. 5. Comparison of the experimental (dots) and theoretical (lines)results for cs(t = 0) = 106 mmoll- I , 53 mmoll- I and 17 mmoI1- 1•

cA zo 4.2 vol.%.

Fig. 6. Diffusion coefficient of oxygen in water for temperaturesbetween 20 and 80 ·C.

6. Non-stoichiometric experiments

So far the Fe(II)NTA chelate solutions were equimo­lar mixtures of NTA and FeS04' Sada has carried outexperiments with an excess amount of NTA to Fe andfound that the rate of oxidation was not influenced byan excess of NTA.

It is equally of interest to investigate the effect of anexcess of Fe, particularly because Neumann and Lynnreport that NTA catalyses the reaction. If this state-

By regrouping the Eq. (7) over for the time course ofthe Fe(lI) concentration the following equation results:

4PA·t (c~-c~)---kR = ~(tO.5 p~.~) (10)

CB - CB

For each series experiments an average for k« is deter­mined (Table 2).

The temperature dependence of kR can be describedin the form of an Arrhenius equation:

kR

= A 'e(-Ea!'1l n

In A and EJ9i can be obtained from linear regressionof the experimental values in a In (kR ) over I/TV plot.

The regression yields an activation energy - Ea =+ 33.10 KJ mol -I and the factor A = 9265 (mmol1- I) - o.s S - I. It is now possible to predict the value ofkR for the temperature range between 25°C and 60 "C,

0.0900.0519

60

0.0700.0387

4025

0.0600.0137

Table 2Values of kR at 25. 40°C and 60 °C

5. Influence of the temperature

development of the Fe(lI) concentration calculatedfrom the solution of the integral. The correlation issatisfactory.

A series of experiments at 40°C and 60 °C werecarried out to examine the influence of the temperatureon the oxidation kinetics of Fe(II). The experimentalset-up remained unchanged. The temperature was im­posed by a heating bath. The temperature not onlyaffects the value of kR, which so far was only deter­mined for 25°C, but also the mass transfer coefficients,the Henry coefficients etc. The Henry coefficients andthe mass transfer coefficient were adapted to the in­creased temperatures.

The mass transfer coefficient follows the temperaturedependence of the diffusion coefficient of oxygen inwater and can be considered to be proportional to thesquare root of the diffusion coefficient.

PL-AThe dependence of the diffusion coefficient on thetemperature is proportional to the temperature depen­dence of the viscosity of water .

D(n = D(2S oC)· P(2S oC)· T (8)298 K· Pm

The dependence of the diffusion coefficient on tempera­ture is shown in Fig. 6. It follows for the temperaturedependence of the mass transfer coefficient that:

DL(n (9)PL(n'a = PL(25 °C).a. DL(250C)

In the case of a known mass transfer and Henrycoefficient for the temperature in question the value forkR can be determined for each experimental value of aseries of experiments.

D. van Velzen et al. / Chemical Engineering and Processing 35 (1996) 443-449

Fig. 7. Development of the Fe(II) concentration with excess Fe(I1).

7. Discussion

A comparison of the results of the present study withthose of Sada et aI. indicates that the values for PLa inboth studies are considerably different. Sada adoptsPLa = 0.036 s - I at 50 °C, whereas in the present studyPLa =0.080 s - I was found. This difference is probablycaused by different reactor dimensions and gas spargersystems. The overall oxidation rate in the experimentsof Sada is found to be ro = 0.33 mmol (1 min) - I forPA=0.05 bar and Co = 5 mmol 1- 1. In the present setup under the same conditions ro is found to be 0.49

oxygenFe(I1)liquid phaseat 11=0

mmol (1 min) - 1. Sada reports a value of kR = 0.058(mmol 1-1) -0.5 s -1 for Fe(II)NTA at 50 °C, whereaswe find k«= 0.041 (mmol I - I) - 0.5 S- I under the sameconditions.

It must be noted that the calculation of kinetic rateconstants for systems where the overall mass transfer ismainly diffusion controlled, always implies the determi­nation of the two contributors (kinetics and diffusion).The contribution of kinetics is determined as the differ­ence between the total rate and the contribution ofdiffusion. The applied numerical value of PLa has thusa very strong influence on the calculated reaction rates.As an example, in the case of Sada a decrease of 5% inPLa would give rise to an increase of k« of 16%. This isdue to the fact that in this set-up the reaction is forabout 75% diffusion controlled. The same effect ispresent to a smaller extent to our work because here thecontribution of diffusion and kinetics are roughlyequal. In the lights of the foregoing, it can be concludedthat the agreement between the results of Sada and thiswork is satisfactory.

ABL1

o

8. Nomenclature

c concentration (mmol l"")CA oxygen concentration in the (mmol l")

bulk of the liquidc~ oxygen concentration at the (mmol I" ')

interfaceDA diffusion coefficient (em" S-I)-t; activation energy (KJ mol-I)Err relative error (%)H Henry's coefficient (bar.l.mol- I)

j number of series IkR reaction rate constant {(mr:IOI) - 0.5S- I)n exponentm exponent

P partial pressure (bar)r reaction rate (unol l") S-I)9l gas constant (8.314 J (mol

K)-I)

1 time (s)T temperature (K)NTA Nitrilotriacetic acidPLa volumetric, overall mass (1 S-I)

transfer coefficient

~ viscosity of water (cP)

8.1. Indices

252010 15t[minl

• ImmolANTA. concentration ~ 3mmolA

• 5mmolA• 7mmolll

5

Oxygen 4.2 vol%Gasflow rate 2041/h

448

10

8

C Fe(II) 6[mmol/ll

4

2

00

ment were true, excess Fe(II) should also be rapidlyoxidised to Fe(III).

A series of experiments was carried out at 25 °C withsolutions of 10 mmol 1- I Fe(II) and 1-7 mmol 1- INTA. The results are shown in Fig. 7.

All curves tend to flatten off to a minimum Fe(II)concentration where no further oxidation occurs. Thisconcentration coincides with the Fe(II) in excess overNTA. It follows that the oxidation rate depends pre­dominantly on the concentration of the unconvertedFe(II)NTA chelate complex.

The explanation is that as soon as all Fe(II)NTA hasbeen converted to Fe(III) NTA the available NTA isused for the complexation of the ferric iron. ExcessFe(II) remains in the uncomplexed state. The oxidationrate of uncomplexed ferrous iron is negligible in com­parison to the oxidation rate of the Fe(II)NTA com­plex.

The conclusion is that the oxidation of Fe(II) is notcatalysed by the presence of NTA. The reaction mecha­nism consists in the conversion of a relatively unstablecomplex (Fe(II)NTA) into a stable one (Fe(III)NTA).

D. van Velzen et al. / Chemical Engineering and Processing 35 (/996) 443-449 449

References

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[2] W. Weisweiler, B. Retzlaff and L. Raible, Simultanabsorptionvon Schwefeldioxid und Stickstoffmonoxid in wassrigen Losun­gen von Eisenchelat-Komplexen, Chem. Eng. Process, 18 (1984)85-92.

[3] W. Weisweiler, R. Blumhofer and T. Westermann, Absorptionof nitrogen monoxide in aqueous solutions containing sulfiteand transition-metal chelates such as Fe(II)EDTA, Fe(II)NTA,Co(II)Trien and Co(II)Tetren, Chem. Eng. Process, 20 (1986)155-66.

[4] S. Tsai, S. Bedell, L. Kirby and D. Zabcik, Field evaluation ofnitric oxide abatement with ferrous chelates, EnvironmentalProgr., 8 (2) (1989) 127-129.

[5] D. van Velzen, H. Langenkamp, D. Papameletiou and H.Nymoen, Verfahren zur Herstellung von Hydroxylamin ausNO x enthaltenden Abgasen, Brevet d'invention Luxembourgeois,Nr. 88 021 17 May 1993.

[6] H. Nymoen, D. van Velzen and H. Langenkamp, Absorption

of NO in aqueous solutions of Fe(II)EDTA: determination ofthe equilibrium constant, Chem. Eng. Process, 32 (1993) 9-12.

[7] J. Hofele, Absorption von Stickstoffmonoxid in waJ3rigen Lo­sungen von Eisen(II)- Chelatkomplexen und deren Regenera­tion durch Desorption und Elektrolyse, Dissertation, TVKarlsruhe, July 1995, EVR Report 16349 DE.

[8] S.S. Jones and FA Long, Complex ions from iron and ethylen­diamintetraacetate: general properties and radioactive ex­change, J. Phys. Chem., 56 (1952) 25-33.

[9] H. Hasui, H. Osuo, H. Ohmichi, Y. Fukuzyu and H. Tarui,Studies of absorption rate of nitrogen monoxide inFe(II)EDTA solution and oxidation rate of this solution (Ab­stract), Nippon Kagaku Kaishi, 3 (1978) 447-455.

[10] D.W. Neumann and S. Lynn, Oxidative absorption of H2Sand O2 by iron chelate solutions, AIChE J., 30 (I) (1984)63-69.

[II] E. Sada, H. Kumazawa, H. Machida, Oxidation kinetics ofFe(II)EDTA and Fe(II)NTA chelates by dissolved oxygen, Ind.Eng. Chem. Res., 26 (1987) 1468-72.

[12] W. Zeise, Zur Reaktionstechnik der simultanen Absorption vonS02 und NO x in wassrigen Eisen- EDTA- Losungen in einerFiillkorperkolonne im Technikumsmallstab, Dissertation,RWTH Aachen, 1989.