Ind. Eng. Chem. Res. 1994,33, 387-394
387
Kinetics of in Situ Degradation of Formaldehyde with Electrogenerated Hydrogen Peroxide Jing-Shan Do' and Chin-Pin Chen Department of Chemical Engineering, Tunghai University, Taichung, Taiwan 40704, Republic of China
In the in situ degradation of formaldehyde with electrogenerated hydrogen peroxide, the cathodic reduction of oxygen on the anodization graphite was first order in the concentration of oxygen dissolved in the aqueous phase. The oxidation of formaldehyde with electrogenerated hydrogen peroxide was second order in formaldehyde and first order in hydrogen peroxide. The activation energy was evaluated as 36 kJ mol-'. A model calculation of the in situ degradation of formaldehyde with electrogenerated hydrogen peroxide correlated well with experimental results. The effects of temperature and initial concentration of formaldehyde on the induction period of the oxidative degradation of formaldehyde with hydrogen peroxide electrogenerated in situ were explored. Introduction The concentration of formaldehyde in waste water must be diminished with some pretreatment before biological treatment because the normal microorganisms die because of the action of the high concentration of formaldehyde therein (Junkermann, 1978). Formaldehyde is effectively degraded with electrogenerated hydrogen peroxide (Do and Chen, 1993a,b). The concentration of formaldehyde is diminished with in situ electrogenerated hydrogen peroxide from 1000 to 2 ppm (Do and Chen, 1993a). The degradation fraction and current efficiency were 99.9 % and 81.1 % ,respectively. The limiting current density of cathodic reduction of oxygen for production of hydrogen peroxide increased by half when the graphite was treated with an anodic current (Do and Chen, 1993b). The mechanisms and kinetics of the cathodic reduction of oxygen to produce hydrogen peroxide with various cathodic materials have been extensively investigated (Morcos and Yeager, 1970;Lovrecek et al., 1983;Paliteiro et al., 1987; Taylor and Humffray, 1975a,b;Yeager, 1984). With pyrolytic graphite and glassy carbon as cathodes, the reaction orders of the cathodic reduction of oxygen in alkaline and acidic solutions were unity (Morcos and Yeager, 1970; Taylor and Humffray, 1975a,b). However, the reaction rate constant and kinetics of cathodic reduction of oxygen to produce peroxide on the anodization graphite are unclear. The reaction orders of the oxidation of formaldehydewith hydrogen peroxide in acidic solution were 2/3 and 1 with respect to the concentrations of formaldehyde and hydrogen peroxide, respectively (Satterfield and Case, 1954). The total reaction order of oxidation of formaldehyde with hydrogen peroxide in acidic solution was 2 (Nikolaev and Ignatov, 1983). The kinetics of oxidation of formaldehyde with hydrogen peroxide in alkaline solution, and the reaction order and activation energy, are still unclear. The mechanism of oxidation of formaldehyde with hydrogen peroxide is not reported. For practical applications, investigation of the mechanism and kinetics of cathodic reduction of oxygen to produce hydrogen peroxide and oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide are important. We investigated systematically the mechanism and kinetics of in situ oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide. A theoretical analysis of the reaction system was compared with
experimentalresults. The exchange current density, Tafel slope, and reaction order of oxygen dissolved in aqueous solution were evaluated experimentally. The reaction orders of formaldehyde and hydrogen peroxide, the rate constant, and the activation energy of oxidation of formaldehyde with hydrogen peroxide were determined. Experimental Section Cathodic Reduction of Oxygen. Graphite was polished with fine emery paper and washed with distilled water in an ultrasonic cleaner. The anodization graphite was obtained by anodic treatment of graphite in phosphate buffer solution (pH 6.88) at 1.6 V (vs Ag/AgCl/3M NaCl solution) until 0.566 C cm-2 was passed. Then the anodization graphite was treated by a cathodic current at 1.5 V (vs Ag/AgC1/3M NaCl solution) for 1min. Experiments on cathodic reduction of oxygen dissolved in the aqueous phase were performed in an undivided cell with anodization graphite as working electrode, a platinum wire as counter electrode, and Ag/AgCl/3M NaCl solution as reference electrode. The current-potential relationships of cathodic reduction of oxygen were obtained by the steady-state method. All potentials are specified to the reference electrode, Ag/AgC1/3M NaCl aqueous solution. Oxidation of Formaldehyde with Hydrogen Peroxide. Formaldehyde was oxidized with hydrogen peroxide in a cylindrical glass reactor. A magnetic stirrer was used to agitate the reaction solution. The entire reactor system was immersed in a water bath, and the temperature was controlled to within f O . l "C. At the beginning of a run, the desired concentration of formaldehyde in aqueous solution (200 mL) was fed into the reactor, which was kept at the desired pH and temperature. The desired stirring rate was applied to the reaction system. After about 10 min, when system equilibrium was achieved, the required amount of hydrogen peroxide was introduced into the reactor. The reaction time was recorded as soon as hydrogen peroxide was fed. Samples were periodically taken from the reactor by use of a hypodermic syringe. The residual oxidant, H202, in the sample was destroyed by NaHS03 (0.01 M). The concentration of formaldehyde was determined according to light adsorption of colored complex of chromotropic acid-formaldehyde at 575 nm (Altshuller et al., 1961). In Situ Degradation of Formaldehyde with Electrogenerated Hydrogen Peroxide. With anodization
* To whom correspondence should be addressed. o s a s - ~ s a ~ ~ ~ ~ t ~ ~ ~1994 ~ -American o ~ ~ ~Chemical ~ o ~ .Society ~ o ~ o
388 Ind. Eng. Chem. Res., Vol. 33, No. 2, 1994
Scheme 1
reaction rate is expressed as
1111I I l l I l l I I I l111l I l l l l l l l l I Ill11 I l l I l I l l l L 202
+ PHCHO + POA- = PHCOI- + P H l o + A,
.I _ . _ _ _ __________.. ..
re = -d[021aq/dt = k2[(0,)dl
Aqueous
___
.._ _ _._ ._ . _ _._ .______ _._ .._ .. . ~ n t e r f a c e
As eq 1 is in equilibrium, [(02-)a&] is obtained as
Gaseous
(Od,
graphite as cathode, the experimental procedures of in situ oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide were executed as described previously (Do and Chen, 1993a).
(6)
[(OL)a&l= (kl/k-1)[(02)aql
(7)
With substitution of eq 7 into eq 6, the rate of cathodic reduction of oxygen on the anodization graphite is
The rate constants of the electrochemical reaction in eq 1 are expressed as (Bard and Faulkner, 1980)
Theoretical Analysis The reaction pathways of in situ oxidative degradation of formaldehydewith electrogeneratedhydrogen peroxide are described in Scheme 1. Oxygen in the gaseous phase transfers into the aqueous phase and is reduced at the cathode to produce hydrogen peroxide. Hydrogen peroxide transfers from the cathodic surface to the aqueous phase and oxidized formaldehyde. The products of this oxidation reaction are formic acid, water, and hydrogen (Do and Chen, 1993a). Cathodic Reaction of Oxygen Dissolved in the Aqueous Phase to Produce Hydrogen Peroxide. The mechanisms of the cathodic reduction of oxygen dissolved in the aqueous phase on the anodization graphite are described as follows (Anastasijevicet al., 1987;Taylor and Humffray, 1975a,b,c;Lovrecek et al., 1983;Paliteiro et al., 1987). The oxygen dissolved in the aqueous phase obtains one electron and adsorbs on the cathodic surface.
in which k10 and 12-10 are the rate constants when the cathodic potential is zero, a is the charge-transfer coefficient, F is the Faraday constant, and E is the cathodic potential. Then
re = k , ~XP(-FE/RT)[(O,)~~I
(11)
where
As E = Ew
+ 7, eq 11 is expressed as
ki
where The rate-determining step is the migration of (02-)a& to the active site on the anodization graphite (Taylor and Humffray, 1975a,b,c;Lovrecek et al., 1983;Paliteiro et al., 1987),
k,' = k, exp(-FE,/RT)
and 7 is the overpotential of the cathode. Therefore the current density of cathodic reduction of oxygen dissolved in the aqueous phase is
i = 2Fk,' The 02-adsorbed on the active site reacts with a water molecule to form a free radical.
The free radical obtains a further electron from the cathodic surface to form an anion that desorbs from the surface into the aqueous solution.
(14)
e~p(-FdRT)l'(O,)~~l
(15)
This relation indicates that the cathodic reduction of oxygen on the anode is first order with respect to oxygen dissolved in the aqueous phase. Oxidation of Formaldehyde with Hydrogen Peroxide in the Aqueous Phase. The oxidation of formaldehyde with hydrogen peroxide occurs in alkaline solution. When hydrogen peroxide is dissolved in the alkaline solution, an equilibrium of hydrogen peroxide is established, kd
H202+ OH- = H0,-+ H20
(16)
ka
ki
(4)
The equilibrium equations of formaldehydein the alkaline solution are (Vaskelis and Norkus, 1991) ki
CH,O
+ H,O * CH,(OH),
(17)
k-7
If the mass-transfer rate of the oxygen from the gaseous phase into the aqueous phase and from the aqueous phase to the cathodic surface is sufficiently large, the reaction rate is controlled by the cathodic reduction of oxygen on the cathode. As the eq 2 is the rate-determining step, the
k0
CH,(OH),
+ OH- =CH20HO-+ H20
(18)
k_s
In the alkaline solution H02- reacts with CH2OHO- to
Ind. Eng. Chem. Rea., Vol. 33, No. 2, 1994 389 form a peroxide (Satterfieldand Case, 1954). The peroxide
O.*
H CH20HO-
I I
ke
+ m-
+ OH-
H-C-0-0-H k-
9
(19)
0-
further reacts with CHzOHO- to form a dimer. A mole of
H
H
I
H-C-0-0-C-H
I
I I
+ H20
(20)
0.15
dimer is degraded to 2 mol of formic acid, H
H
I I
H-C-O-O--C-H
0-
I I
kll
2CHQ-
0.20
0.25
-E, V (vs. Ag/AgC1/3M
+&
(211
0.30
NaC1)
Figure 1. Steadystate polarization curves of the cathodic reduction of oxygen dissolved in the aqueous phase: cathode, anodization graphite; cathodic area = 60.4 cm2;pH = 13; temperature = 46 O C .
0-
is expressed as
If the reaction in eq 19 is the rate-determining step, the reaction rate of the oxidation of formaldehyde with hydrogen peroxide is -d[CH,OI/dt = kg[CH20HO-l[HO~l (22) The concentrations of CH20HO- and HO2- obtained based on eqs 16,17, and 18 are substituted into eq 22. fhl=
k$($(,K,ICHzOlCH2021IOH-12 (23) where KS = ks/k_s, K7 = k7/k-7, and K8 = k$ka. When the pH value is kept at constant, the reaction rate is expressed as rh1=
= kh1[CHzOl [H2021
= kM[CH2012[H2021
(29) in which kh2 = kll'[OH-l2. The results reveal that the reaction orders of oxidation of formaldehydewith hydrogen peroxide are 2 and 1with respect to formaldehyde and hydrogen peroxide, respectively. Mass Balance of Oxidation of Formaldehyde with Electrogenerated Hydrogen Peroxide. As shown in Scheme 1, the rate of generation of hydrogen peroxide from the cathodic surface is expressed in eq 15. The rate of oxidation of formaldehyde results in consumption of hydrogen peroxide. Therefore, the mass balance of hydrogen peroxide in the bulk solution is rh2
(24) where k h l = kgK&7Ks[OH-12. The result indicates that the reaction orders of formaldehyde and hydrogen peroxide are all unity. If the reaction in eq 21 is the rate-determining step, the reaction rate of the oxidation of formaldehyde with hydrogen peroxide is expressed as
d[H,OJ/dt = GA/2FV) - 0.5kh[CH201m[H20.J" (30) where A and V are the area of cathodic surface and the volume of solution, kh, m, and n are the rate constant and the reaction orders of formaldehyde and hydrogen peroxide. The concentration of formaldehyde in the bulk solution is expressed as
rh2 = -d[CHzOl/dt = 2kii[Dl in which D is
d[CHzOl/dt - ~ ~ [ C H Z ~ ] ~ [ H ~ O(31) ~]" The concentration of formaldehyde and hydrogen peroxide in the aqueous phase for kinetic control are evaluated by solving eqs 15, 30, and 31 simultaneously when the kh, m, and n are obtained experimentally and the cathodic reduction of oxygen is kept at potentiostate. When a constant current density is applied to the cathodic reduction of oxygen to produce hydrogen peroxide, the concentrations of formaldehyde and electrogenerated hydrogen peroxide are evaluated by the solving eqs 30 and 31 simultaneously.
T'hl
H
I
H-C-0-0-C-H
I
0-
(25)
H
I I
0-
According to the equilibrium equations, 16-20, the concentration of the dimer is obtained as
[Dl = K6~K,~2~K~~2K$(10~CH2012~Hz021 [OH-12 (26) Results and Discussion where KO= kolk-9 and KIO= klo/k-lo. Substituting eq 26 Kinetics of Cathodic Reduction of Oxygen. A. into eq 26 gives Effect of Concentration of Oxygen Dissolved in the Aqueous Phase. The steady-state polarization curves of rh2 = kl,'[CH2012~H,021[OH-12 (27) cathodic reduction of oxygen on the anodization graphite in which are shown in Figure 1. The current density of cathodic reduction of oxygen increased when the cathodic potential kl,' = 2kllK6(K7~2~K~2K&,, (28) increased. When the concentration of oxygen dissolved in the aqueous phase was increased from 0.144 to 0.678 If the pH of the reaction solution is kept constant, eq 27
390 Ind. Eng. Chem. Res., Vol. 33, No. 2, 1994 -9.0
Temp., K
-9.5 N
'E0 4 -10.0
.c 4 -10.5
-11.0 -9.0
/ -8.5
slope=0.99*0.02
-8.0
-7.5
-0.15 -5.25
-7.0
-4.75
-4.25
-3.75
-3.25
Log j , A cm+ Figure 2. Effect of concentration of oxygen dissolved in the aqueous phase on the current density of cathodic reduction oxygen: cathode, anodizationgraphite; cathodic area = 50.4cm2;pH = 13;temperature = 45 OC; cathodic potential = 0.22 V.
mM, the current density of the cathodic reduction of oxygen increased from 0.022 to 0.095 mA cm-2 when the cathodic potential was kept at 0.22 V. A plot of the logarithmiccurrent density of cathodic reduction of oxygen against the logarithmic concentration of oxygen dissolved in the aqueuous phase yielded a straight line with a slope of 0.99 f 0.02, which was the electrochemicalreaction order of oxygen as shown in Figure 2. The experimental results correlated well with eq 15. This agreement indicates that the derived mechanism described in the theoretical analysis is justified. B. Effect of Temperature. Rearranging eq 15, we obtained j = j , exp(-FqlRT)
(32)
Figure 3. Tafel plot for various temperatures: cathode, anodization graphite; cathodic area = 50.4 cm2;pH = 13;02 sparging rate = 2 cm3 8-1. Table 1. Effect of Temperature on the Cathodic Reduction of Oxmen. -temp, K [0zlaS,mM j ~ ,p, A cm-2 '.k X 1@ cm s-l Tafel slope 1.62 68 288 1.160 3.62 298 0.913 3.85 2.19 68 3.45 64 308 0.753 5.02 6.31 4.81 64 318 0.678 a Cathode, anodization graphite; cathodic area = 50.4cm2;pH = 13;anode, Pt wire.
-9.5
-
io= 2FIz,'[(02),,l
(33)
The exchange current density j o was constant when the concentration of oxygen dissolved in the aqueous phase and the temperature were fixed. The Tafel curves varied with temperatures, and a linear relationship for each Tafel curve was found in the range of cathodic potential between 0.05 and 0.10 V as shown in Figure 3. The exchange current densityj, was evaluated by extrapolation of the linear relationship to B = 0. The saturated concentration of oxygen dissolvedin the aqueous phase decreased from 1.16to 0.678 mM, and the exchange current density of cathodic reduction of oxygen increased from 3.62 to 6.31 pA cm-2 when the temperature was increased from 288 to 318 K (Table 1). On substitution of the exchange current density and the concentration of oxygen dissolved in the aqueous phase into eq 33, the values of rate constant, I+,', were evaluated as shown in Table 1. A straight line was obtained from a plot of the logarithmic rate constant against reciprocal of temperature as shown in Figure 4. The slope of this line was -3391 f 191,which corresponded to an activation energy of 28.2 kJ mol-'. The frequency factor was evaluated as 2.04 cm s-l. The Tafel slopes of linear sections of q-log j curves varied from -68 to -64 mV when the temperature was increased from 288 to 318 K (Table 1). According to eq 32, the
\
slope=-3391* 191
v)
E0
where
-10.0
I
-10.5
J c
4
-11.0
-11.5
3.00
3.25
3.50
3.75
1 / ~ x 1 0 ~ I, / K Figure 4. Effect of temperature on the rate constant of cathodic reduction of oxygen: cathode, anodization graphite, cathodic area = 50.4 cm2;pH = 13;0 2 sparging rate = 2 cms d.
theoretical value of the Tafel slope was -59 mV which agreed satisfactorily with the experimental results. Kinetics of Oxidation of Formaldehyde with Hydrogen Peroxide. A. Effect of Concentration of Formaldehyde. The initial rate of oxidation of formaldehyde with hydrogen peroxide increased from 0.52 to 1.44 mM min-1 when the concentration of formaldehyde was increased from 13.2 to 22.4 mM as shown in Figure 5. A plot of the logarithm of the initial rate of oxidation of formaldehyde against logarithm of the concentration
Ind. Eng. Chem. Res., Vol. 33, No. 2, 1994 391 -6.0
I
-4.0
I
3
-6.5
4 x
- -7.0
I
i
c
d k
v
d
E
-5.0
d k
//
W
c I 4
-5.5
-6.0
-8.0 -5.0
1
slope= l . O l * 0.07
2
J
-7.5
I
I
-4.5
'i
/
slope= 1.94*0.12
2
J
7
I
-4.5
-4.0
-5.0
-3.5
Ln [CH20],M
-4.5
-4.0
Ln
[H202]I
-3.0
-3.5
M
Figure 5. Effect of concentration of formaldehyde on the initial reaction rate of oxidation of formaldehyde with hydrogen peroxide: pH = 13;temperature = 5 OC; [HzOz] = 0.03 M; stirring rate = 600 rpm.
Figure 6. Effectof concentrationof hydrogenperoxide on the initial reaction rate of oxidation of formaldehyde with hydrogen peroxide: pH = 13;temperature = 5 "C; [HCHO] = 0.067 M, stirring rate = 600 rpm.
of formaldehyde yielded a straight line of slope 1.94 f 0.12. This value was approximately equal to 2 and correlated well with the case when eq 21 was the ratelimiting step in the theoretical section. The 2/3order with respect to the formaldehyde in the acidic solution was obtained in the literature (Satterfield and Case, 1954).In acidicsolution,the reaction of formaldehyde with hydrogen peroxide was probably fairly complex and might be associated in some way with existence of polymers in solution (Satterfield and Case, 1954). However, formaldehyde was completely dissociated toCH2(0H)2in alkaline solution when the concentration of formaldehyde was below 0.5% (Murphy et al., 1989). The mechanism of oxidation of formaldehyde with hydrogen peroxide in alkaline solution was derived in the theoretical section. When eq 21 was the rate-determining step, the reaction order of formaldehyde was 2 and correlated well with the experimental results. B. Effect of Concentration of Hydrogen Peroxide. Experimental results demonstrated that increasing the concentration of hydrogen peroxide from 10.0 to 30.0 mM and keeping the concentration of formaldehyde a t 67.0 mM resulted in an increase of the initial rate of oxidation of formaldehyde from 3.55 to 10.80 mM min-l. The plot of the logarithmic initial rate of oxidation of formaldehyde against the logarithmic concentration of hydrogen peroxide yielded a straight line as shown in Figure 6. The slope of the straight line was 1.01 f 0.07 which was the reaction order of hydrogen peroxide. The reaction order of hydrogen peroxide agreed with our expectation. It also revealed that the proposed mechanisms were reasonable and the rate-determining step was eq 21. The values of m and n in eq 31 were obtained as 2 and 1, respectively. The rate constant k h in eq 31 was equal to kh2 in eq 29. C. Effect of Temperature. Substitution of the concentrations of formaldehyde and hydrogen peroxide into eq 31 yielded the value of rate constant k h of the oxidation of formaldehyde. When the temperature was increased from 278 to 298 K,the initial reaction rate of formaldehyde increased from 1.14 X 10-3to 3.17 X 10-3 M min-' and the rate constant k h increased from 94.8to 264.5 M-2min-1 (Table 2). A plot of the logarithmic rate constant of the oxidation of formaldehyde against the reciprocal of temperature resulted in a straight line as shown in Figure
Table 2. Effect of Temperature on the Reaction Rate and Rate Constant of the Oxidation of Formaldehyde with Hydrogen Peroxide. temp, K 278 283 288 293 298
(h)iX
109,M min-1 1.14 1.38 1.86 2.42 3.17
min-1 94.8 115.0 154.7 201.5 264.5
kh, M-2
a pH = 13;stirring rate = 600 rpm; [HCHO] = 0.02 M; [HzOz] = 0.03 M.
6.0
* I
I
*
5.5
c
\2
slope=-4329%?77
*2
N
E
5.0
\
z
\
d e I4 4.5
4.0
3.00
3.25
3.50
3.75
4.00
1 / ~ ~ 1 0I / ~ K, Figure 7. Effectof temperatureon the rate constant of the oxidation of formaldehyde with hydrogen peroxide: pH = 13;[HzOzl = 0.03 M [HCHO] = 0.02 M stirring rate = 600 rpm.
7. The slope of the straight line was -4329 f 277,which corresponded to an activation energy of 36.0 kJ mol-'. The intercept of the straight line was 20.12 f 0.96,and the pre-exponetial factor was evaluated as 5.47 X 108M-2 min-1. In Situ Oxidative Degradation of Formaldehyde with Electrogenerated Hydrogen Peroxide. Current Efficiency and Concentration Distribution. As shown by the dashed line of Figure 8,the experimental results
392 Ind. Eng. Chem. Res., Vol. 33, No. 2, 1994 100
--
E
:
-
'-
: deeradatlon of H C H O
-1
experimental data
800
E
a a
E
I
c;
2
u2
200fi 1 I
/
'\\
0
200
' \\ $ \ \
0
I
50
0
100
150
, ,,
200
250
;' 20 0 300
0 0 ' 0
I
I
I
100
200
300
Time, min Figure 10. Effect of reaction time on the current efficiency of the oxidative degradation of formaldehyde and the production of hydrogen peroxide: cathode, anodization graphite; cathodic area = 50.4 cm2;[HCHO] = lo00 ppm;current density = 0.75 mA cm-2; pH = 13; temperature = 45 OC; 0 2 sparging rate = 2 cm3 s-1.
*
a
u
: :
O* sparging rate=:! N~ sparging rate=2
em's-' ern's-'
& p o
eR
90
z
.IIi
'O0:
was 99.7 5% when the 0 2 sparging rate was 2 cm3 s-l. The experimental results revealed that the effect of direct reduction of formaldehyde on the degradation of formaldehyde was slight. The current efficiency for production of hydrogen peroxide decreased from 100% to a minimum value of 8.6 % when the period of reaction increased from 0 to 67 min (Figure 10). The current efficiency of the oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide increased from 0 to a maximum value of 91.4%. At the initial stage, most of the current was used to generate the oxidant hydrogen peroxide (02)aq + H20 + 2e
0
200
406
Charge Passed, C Figure 9. Effect of charge passed on the concentration and degradation of formaldehyde: cathode, graphite; cathodic area = 50.4 cm*; [HCHO] = 500 ppm; current density = 0.50 mA cm-2;pH = 13; temperature = 45 O C .
indicated that the concentration of formaldehyde decreased from 1000 to 10 ppm when the period of oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide increased from 0 to 225 min. The theoretical calculations correlated well with the experimental results as illustrated by the solid line of Figure 8. In the comparison between experimental results and theoretical calculations,the concentration of formaldehyde obtained from the experimental results decreased more rapidly than that from the theoretical calculations. Possibly, formaldehyde in the aqueous phase was degraded by oxygen dissolved in the aqueous phase (Do and Chen, 1993a)and some formaldehyde may also have been directly reduced on the cathodic surface in addition to degradation of formaldehyde with electrogenerated hydrogen peroxide. As shown in Figure 9, the degradation of formaldehyde due to the direct reduction of formaldehyde on the cathode in the absence of 0 2 was 11.8% when 300 C was charged into the electrolysis cell. However, the degradation of formaldehyde with electrogenerated hydrogen peroxide
-
HO;
+ OH-
(34) and the current efficiency for production of hydrogen peroxide was larger than that of oxidative degradation of formaldehyde. When the reaction duration was increased, the oxidative degradation of formaldehyde increased due to increased concentration of electrogenerated hydrogen peroxide. 2HCHO + HO;
+ OH-
-
2HCO;
+ H20 + H2
(35)
A t this stage, the most charge passed was used to degrade formaldehyde in the bulk aqueous solution. The current efficiency to produce hydrogen peroxide decreased and the current efficiency for the oxidative degradation of formaldehyde increased. With further increased reaction duration, the decrease of concentration of formaldehyde in the aqueous phase resulted in decreased current efficiency of oxidative degradation of formaldehyde and increased current efficiency of production of hydrogen peroxide (Figure 10). Induction Period. As shown in Figure 8, the reaction system had an induction period during which the concentration of hydrogen peroxide increased rapidly and that of formaldehyde decreased slowly. A t the beginning of the run, the rate of reaction 34 was larger than the rate of 35. The concentration of hydrogen peroxide in the aqueous solution therefore increased in the initial period of a run. After the induction period, the rate of reaction 35 increased and was equal to the reaction rate of eq 34. The system was in a pseudo-steady-state. The concen-
Ind. Eng. Chem. Res., Vol. 33, No. 2, 1994 393 I
1000
I
: conc. of HCHO
l4
E
2
900
d
1: 318 2: 308 3: 298
3 w
1
800
0
..- 5
I
600:-
I
1
I
10
20
30
'0 40
Time, min Figure 11. Effect of reaction time on the concentrations of formaldehyde and hydrogenperoxide: cathode,anodizationgraphite; cathodic area = 50.4 cm2; [HCHO] = lo00 ppm; current density = 0.75mA cm"; pH = 13;0 2 spargingrate= 2 cm3s-1. -, concentration of formaldehyde; - - -: concentration of hydrogen peroxide. Table 3. Effect of Temperature and the Concentration of Formaldehyde on the Induction Period. temp, K 278 288 298 308 318 318 318 318
[HCHOli, ppm lo00 lo00 1000 lo00 lo00 1250 750 500
induction period, min 10.0 7.4 5.1 3.9 2.6 1.8 4.6 7.0
Cathode,anodizationgraphite;cathodicarea = 50.4 cm2;current density = 0.75 mA cm-2; 02 sparing rate = 2 cm3 s-l; pH = 13. 0
tration of formaldehyde decreased significantly, and the concentration of electrogenerated hydrogen peroxide increased slightly. The induction period of oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide increased as the temperature was decreased (Figure 11). When the temperature was decreased from 318 to 278 K, the induction period of the reaction system increased from 2.6 to 10.0 min (Table 3). When the temperature of the reaction system was decreased, the rate constant of oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide in eq 35 decreased. As a constant current was applied, the rate of production of hydrogen peroxide, eq 34, was kept constant when the reaction temperature was decreased. The concentration of electrogenerated hydrogen peroxide in the aqueous solution should be increased to compensate for the decrease in the reaction rate constant to keep the system in the pseudosteady-state when the reaction temperature was decreased. Therefore the induction period increasedwhen the reaction temperature was decreased. As illustrated in Table 3, when the concentration of formaldehyde was increased from 500 to 1250 ppm, the induction period decreased from 7.0 to 1.8 min. The pseudo-steady-state concentration of hydrogen peroxide electrogenerated on the cathodic surface decreased and resulted in the decrease of the induction period when the concentration of formaldehyde was increased.
Conclusions The theoretical calculations of oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide correlated well with experimental data. The reaction order of cathodic reduction of oxygen was 1 with respect to oxygen dissolved in the aqueous phase. The reaction orders of the oxidation of formaldehyde with hydrogen peroxide were 2 and 1 with respect to formaldehyde and hydrogen peroxide, respectively. The activation energies of cathodic reduction of oxygen and oxidation of formaldehyde with hydrogen peroxide were 28.2 and 36.0 kJ mol-', respectively. The induction period of oxidative degradation of formaldehyde with electrogenerated hydrogen peroxide increased when both the temperature and concentration of formaldehyde were decreased. Using electrogenerated hydrogen peroxide as oxidant, formaldehyde was degraded from 1000 to 10 ppm within 225 min. Acknowledgment The support of the National Science Council of the Republic of China (NSC 80-0410-E029-03) and Tunghai University is acknowledged. Nomenclature A = surface area of cathode E = cathodic potential E, = equilibrium cathodic potential F = Faraday's constant = current density J O = exchange current density k = rate constant K = equilibrium constant kO = electrochemical rate constant at E = 0 kill = 2 k l l ~ 1 3 ( ~ 7 ) ~ ( ~ 8 ) ~ ~ & 1 0 k, = kz(klO/k-lO) k,' = k, exp(-FE,/R!l') khi = ksK&7KdOH-12 kh2 = kl1'[OH-l2 r = reaction rate t = time T = temperature V = volume of solution Subscripts aq = aqueous phase ads = adsorption i = initial state Greek Symbols = charge-transfer coefficient 9 = overpotential CY
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Received for review June 1, 1993 Revised manuscript received October 2, 1993 Accepted October 12, 19930 Abstract published in Advance ACS Abstracts, December
1, 1993.
