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Activation energy variation for catalytic oxidation of aqueous SO2
ACTIVATION ENERGY OXIDATION
VARIATION FOR CATALYTIC OF AQUEOUS SO,t
W. PASIUK-BRONIKOWSKA* Institute of Physical Chemistry, Polish Academy
and A. SOKOEOWSKI of Sciences, 01-224 Warszawa, Kasprzaka44/52.
Poland
(Received 12 May 1982; accepted 16 August 1982) oxidation in aqueous solutions catalysed with manganous sulphate was studied to determine Abstrart-SO, temperature dependencies of the reaction rate. The process was carried out at relatively high sulphuric acid (reaction product) concentrations with regard to its application for SO2 removal from waste gases. Variation of the apparent activation energy has been linked with alteration of reaction rate determining steps. INTRODUCTION
The catalytic influence of some transition metal ions on oxidation of SO, absorbed in aqueous solutions has been known since the 19th century. However, there arc at least two important reasons why research on this process is
still
carried
on.
One
is
its
complexity
solution into the reactor, mocouple for indicating mixture. Both the reactor tem were immersed in a within f O.l”C.
and a Pyrex jacketed thertemperature of the reacting and the gas conditioning syswater bath and thermostatted
involving
laborious studies on the process mechanism and another one is its practical importance for SO1 emission control. To develop such a technology sufficiently safe data are needed, particularly the optimum catalyst, its concentration, relative concentrations of reagents, reaction temperature etc. In this work we attempted to determine experimentally the activation energy of oxidation of SO, absorbed into the MnSO. aqueous solution hoping to gain additional information as to the reliability of the reaction model proposed prcviously[l]. Our intention was also to explain discrepancies in the scarce literature on the intlucncc of temperature upon the process. Some authors reported relatively high values of activation energy (Hoathcr et 01.[2]-27.3 kcallmol ? 4%, Huss[3]--19.8 + 0.7 kcallmol), whereas others indicated very low ones (Tarbutton et a/.[41 did not observe any measurable effect of temperature. the value evaluated from the data of Copson et al. [5] is lower than 2 kcal/mol). EXPERIMENTAL
Apparatus Experiments on SO2 oxidation were conducted in a semibatch foam reactor with constant flow of gaseous reactants, previously passed through a mixer and a humidifier. We have chosen this type of gas-liquid contact as it is directly transferable to industrial scale. A vertical Pyrex tube was used as the reactor (Fig. I), supplied at the bottom with the fine glass-frit as a gas distributor and, at the top with the PTFE lid with through-pipes. The pipes were coupled with a reflux corklenscr discharging after-reaction gases, tube for liquid sampling, separatory funnel applied for introducing initial or excessive (taken through the sampling tube) ‘Author to whom correspondence should be addressed.
tPart of this work was presented at the European Conference
of the Federation of European Chemical Societies: Chemical Pathways in the Environment, Palaiseau, France 1980.
ikiaten’als Sulphur dioxide. The supply of sulphur dioxide was from a technical cylinder placed outside through an intermediate steel bottle (2 I.) placed in the vicinity of the reactor and periodically loaded by distillation of liquid SO* from the outer cylinder. The gas was also cleaned by passing it through a silica gel column and then through a fine porous glass plate. Air. Molecular oxygen from air was used as an oxidizing agent. Air was sucked from the outside of the laboratory with a diaphragm pump and next passed through a silica gel column, cloth titer and fine porous plate. Manganous s&hate. The analytical grade reagent was from PPH POCH, Gliwicc. It was applied without any further purification. Water. All solutions were prepared with redistilled water. Procedure As soon as the bath temperature was fixed at the desired level the flow of air was turned on and then the reactor filled with the catalyst aqueous solution of known volume and Mn concentration. Simultaneously with the solution the proper stream of SO, was introduced into the flowing air and hence the oxidation was under way. From the start of supplying SOI a run was timed and successive liquid samples were withdrawn. Only 1 ml liquid portions were needed for analysis and the excess of the withdrawn liquid was returned to the reactor. The temperature of the reflux condenser was adjusted according to the programme worked out on the basis of preliminary experiments so as to keep the possibly constant volume of the reacting liquid, independently of the increasing H2S04 concentration with the progress of oxidation. Samples were analysed for H2S04 to get information on the extent of reaction and for Mn to allow correction for changes in the reacting solution volume. Alkalimetric
122
W.
PASIUK-BRONIKOWSKA and
A.
SOKOLOWSKI
Fig. 1. Schematic drawing of experimental apparatus: 1, pyrex tube 0.07m i.d.; 2, gas distributor/glass-friitl; 3, PTFE lid: 4, reflux condenser: 5, separatory funnel: 6, sampling tube; 7, thermocouple; 8, humidifier; 9, mixer of gaseous reagents; 10, rotameter; 11, Ueometer; 12, fine porous plate; 13, cloth filter; 14, silica gel column; 15, diaphragm pump. titration with metal masking and calorimetric determination in the presence of formaldoxime were employed respectively. Below fundamental parameters of experiments are specified: MuSO concentration SO? concentration flow of gaseous mixture liquid volume temperature
1.2 x lo-‘-0.18 mol/dm3 0.5-3 vol.% 0.1 l-O.44 dm’/s 0.134.25 dm’ 14.2-43.3’C.
RILWLTSANDDISCUSSION The rate of sulphuric acid production which is equal to the rate of SOz oxidation can be obtained from the equation r = (dnJdt)/
V = dc,/dt
- (dcJdt)(cJcM)
(1)
where n. = n.(t) and V= V(t) were not measured while c, = c,(t) and CM = c&t) were experimentally determined concentrations of sulphuric acid and manganese, respectively. The data points were expressed as higher order polynomial functions of time using the least squares method. Exemplary fitted curves are shown in Pi. 2. In view of previous reports[l-3.51 as well as of observations made in this work sulphuric acid causes significant retarding effect on the rate of its production. Therefore to examine the reaction sensitivity to temperature one should compare reaction rates at diRerent temperatures but alike acid concentrations or reaction extents when starting with an aqueous solution containing no acid. Temperature dependencies for SO2 oxidation at MnSOa concentration 5 x lo-” mol/dm3 and various
5
10
15 tno-T
Fii.
20
25
30
S
2. Sulpburic acid and manganese concentration 14.2’C; 0, 25.1”C and Cl, 34.8”C.
va time;
A,
reaction extents are plotted in Fig. 3. It shows results of experiments given as In r vs l/T instead of the typical Arrhenius plot, as the process is complex and the form of its rate equation may he still in question. Under the condition that experiments were performed at constant concentrations of both substrates the value of the apparent activation energy found from Fig. 3 did not differ from that determined from the classic Arrhenius equation. Data points reported in Fig. 3 are arranged in two regions. one of variable activation energy which
Activation energy variation for catalytic oxidation of aqueous SOI
I n
O,l
0.2 “_
-t
123
a3 .
Ob
I
dnP/S
<
Fi.
f
1
k
“’ L
I
32
, 3.3
3.L iO?T,
K-’
3.5
4. Dependence of the process rate upon gas flow rate for cM= 5 x IOFmolldm’and temperature 32.PC.
more probable. For this purpose experiments were carried out at all parameters constant but the gas Bow rate which was varied to change hydrodynamic conditions determining transfer of gaseous reactants. Figure 4 shows that at the gas flow rate 0.236 dm’ls applied in the main set of our experiments the influence of hydrodynamic conditions was not detectable. Rates of SO2 oxidation plotted in Fii. 4 are the maximum ones attained during runs at 32.7”C (found by extrapolation to c. = 0). These results allow more firmly to ascertain that the simultaneous absorption of oxygen and sulphur dioxide accompanied by the catalytic oxidation of the latter goes in the kinetic regime throughout tbe conditions described in Fii. 3 up to the temperature as in Pii. 4. However, further increase in the temperature might partly bring dilfusional limits on the examined oxidation process. The effect of temperature on the rate of SO2 oxidation at the relatively bigb MnSO., concentration (up to 0.18 mol/dm3) was rather weak. It is illustrated in Fig. 3 (dotted line) as derived from initial process rates of experiments with upper catalyst concentrations (see diffusional regime in Fig. 5). In this case the liquid side resistance of the oxygen transport was the rate -controlling phenomenon. Nevertheless, the foam reactor is not convenient for studying the process in the diffusional regime. It is mainly because the area of gas-liquid interface, especially foam height, is dependent upon the physical properties of the reacting mixture which change with the reaction going on. For SOa oxidation catalysed with MnSOd in aqueous solutions Pasiuk-Bronikowska and Bronikowski[l] have reported that the SO, absorption rate may be given by the empirical relation:
I
Pi. 3. Temperature dependencies for SO1 oxidation at c, = 5 x lo-’mol/dm’(exemplary points: + , c, = 0 mol/d&; A, c, = 0.408moI/dd; 0, E. = 0.816 mol/dm’). Values of activation energy at higher temperatures (left side of the diagram) are: 1.03. 2.75, 4.32. 6.10. 8.83. 10.70kcai/mol respectively with c. increasing.
decreases by several times with lowering sulphuric acid concentration and another one of almost constant activation energy 24 f 2.4 kcal/mol. The latter is extended to higher concentrations of sulpburic acid and lower temperatures. To get further information on the character of the regions it was necessary to answer whether they were diffusional or kinetic. Therefore fluxes of gaseous substrates (sulphur dioxide and oxygen) from the gas phase into the liquid bulk were calculated on the basis of previously found rates of oxidation as well as of the estimated mass transfer parameters k, = 5.5 x IO-’ m/s and (I = 14.5 x 10’rn-‘, and their values compared. The mass transfer parameters were determined from the correlations suggested by Gestrich et af.161. The comparison indicated that tbe slow reaction regime could be applied to both regions. The calculations also disclosed that with the additional increase in the rate of oxidation one might expect the appearance of the liquid side resistance first due to oxygen and not to sulphur dioxide. Oxygen diffusional limitation calculated on this basis is shown in Fii. 3 (dashed lime). As the estimation of kr and a was rather rough it seemed reasonable to verify practically the kinetic regime, especially for the variable activation energy region. In this instance the diffusional limitation wa*
llr= AdeM
+ B(c,/c,‘)
(2)
llr = A&,
+ B(c,/cM2)
(31
or by
depending on relative rates of appropriate step reactions. Here, the equations have been written in a converted form which allows to visualize that by examining a run
W. PASIUK-BRONIKOWSKAadd A. SOKOLOWSKI
124
where cm = [Mn”]([Mn3’] was negligible in comparison with the concentration of MnS04 throughout experiments), k, is a composite rate constant being a product of appropriate step reaction and equilibrium constants, and K’ and K” are equilibrium constants of particular step reactions. Comparison of eqn (3) with eqn (4) leads to A,= K’lk, and B = Wk,. When co= const which may be a rough approximation for a run eqn (4) becomes:
where k, = kc0 and K, = K”(co/cM2). As results from the equations contribution of the oxygen solubility temperature effect to the empirical value of the apparent activation energy should be considered. This effect is expected to be excluded when K,c, B K’. In this case lo-’
IO-'
1w2 cM
,
W, = (kJK”)cM’.
(6)
mol/dm’
Fii. 5. The influence of c,,, on the slope of the line l/r = f(c,) (eqns 2 or 3) disclosing diffusional limitation. as l/r vs c, the regime of gas absorption may be easily recognized. For both cases (eqns 2 and 3) when the kinetic regime (slow reaction) takes place the slope of a line should be inversely proportional to the squared M&O4 concentration (S = B/ca) provided all relevant parameters are kept constant. Similarly the intercept I = A Jc,’ (eqn 2) can be based on for the reaction going in the 0, independent region. The regularity vanishes in the diiusional regime. This is demonstrated in Fig. 5. It gives the slopes calculated for several runs performed at different values of cIK and temperatures 20.0”, 32.7’ and 40.0°C, plotted against c,,, (log-log coordinates). The data points follow the expected relationship up to a certain value of CM,being lower for the higher reaction temperature. Further increase of the MnS04 concentration indicates transition from the kinetic to the diiusional regime for which the above relationship ceases to hold. The results conlirm that conditions for all experiments shown in Fig. 3 below the dotted line corresponded to those of the kinetic regime. Even at the temperature 4O.O“Cthe intbtence of diffusion appeared only when cM >6~ lo-’ mol/dm” whereas experiments in Fii. 3 were at C~ = 5 x lo-“mol/dm3. Hence it can be concluded that the unusual effect of temperature observed in Fig. 3 reveals mainly the complex reaction mechanism. As the plot is curved downward this is not the case of parallel pathways but of reaction successive steps. According to Pasiuk-Bronikowska et a/.[11 who proposed the reaction model for oxidation of SOI catalysed with MnSOd the theoretical rate equation for the set of parameters applied in this work may be expressed as:
Plotting experimental data as rc, against c, the approached values of (t/K’9cM2 = const could be found for several temperatures in the lower range and hence the single apparent activation energy as given in Fig. 6 (a and b-l). Values of the coefficient (&,/K”)c,’ were also calculated from eqn (4) takine k&K’ as the initial rate of oxidation at c. = 0 (see b-2 in Fig. 7). Thus obtained values for the activation energy are 20.2 -(b-l) and 23.4 kcal/mol (b-2) indicating discrepancy caused by the way of data treatment. The results bear a slight compensating effect with respect to the second term of denominator in eqn (4) as co decreases with the rise of c.. Huss[31 who supplied the best evidence for his experiments gave the value of activation energy found at c0=7.S x lo-*mol/dm3, c, = 0 and c, = 3.31 x lo-’ mol/dm3 for temperatures between 25 and 38°C. In view of previous considerations on reaction mechanism[l] the conditions were fulfilled for the kinetics described by:
r = kocM2.
(7)
Therefore values of activation energy found in this work and reported by Huss should not necessarily be equal. CONCLUSIONS
Variation of activation energy for the reaction of SO? oxidation catalysed with Mn has been proved experimentally. To cbstinguish between the tirst order reaction and diiusion controlled kinetics with respect to oxygen the influence of other reagents (MnS04 and H2S04) in the complex reaction was successfully examined. Attempts were made to link such a behaviour with the reaction mechanism, which allowed to settle that SO* oxidation kinetics may or may not significantly depend on temperature according to the reaction order with respect to oxygen. With the order increasing (from 0 up to 1) the value of apparent activation energy diminishes from 19.8 kO.7 kcallmol in the absence of HJO, or 24%
Activation energy variation for catalytic oxidation of aqueous SO*
a
Fig. 6. Determination
24kcal/mol at relatively high to about 1 kcal/mol. The latter tors at high Mn concentrations overcome the energy-related acid contamination.
of apparent activation energy basing on eqn (6) (a and b-l) or eqn (4) (b-2).
concentrations of H2S04 is attained in foam reacencountered in practice to problem of environment
NOTATION
specific interfacial area with respect to gas-liquid volume, m-’ kinetic constants in experimental rate equations, mol s/dm’ and s respectively (eqns 3 and 4) kinetic constant in experimental rate equations (3 and 4), s manganous sulphate concentration, molldm3 oxygen concentration (in liquid), mol/dm3 sulphuric acid concentration, mol/dm’ apparent activation energy of the reaction (process), kcallmol intercept for I/r = f(c.) (oxygen independent region), dm3 s/m01 liquid side mass transfer
coefficient,
m/s
complex
rate constant
complex equilibrium constants moles of sulphuric acid, mol rate of the reaction, molldm’ s slope for i/r = f(c,), s/mol* dm6 time, s temperature of the reaction (process), “C absolute temperature, K volume of reacting liquid, dm3 gas flow-rate, m”ls
REFERENCF.S
Cl1 Pasiuk-Brouikowsks
Sci. 1981 36 215.
W. and Bronikowski T.. C/rem. Ensare - _
Dl Ho&her R. C. and Goodeve 30 11‘49.
C. F., Trans. Faraday
Sot. 1934
Huss A. Jr., Ph.D. Thesis. Urbana. Illinois 1978. E:; Tarbutton G., Driskell I. C., Jones T. hf., Gray F. J. and Smith C. M., bd. Engng Chem. 1957 49 392. PI Copson R. L. and Payne J. W., Ind. Engng Chem. 1933 25 9m ___. WI G&rich W., Esenwein H. and Kmuss W.. ht. Chem. Engng 1978 18 38.
VARIATION FOR CATALYTIC OF AQUEOUS SO,t
W. PASIUK-BRONIKOWSKA* Institute of Physical Chemistry, Polish Academy
and A. SOKOEOWSKI of Sciences, 01-224 Warszawa, Kasprzaka44/52.
Poland
(Received 12 May 1982; accepted 16 August 1982) oxidation in aqueous solutions catalysed with manganous sulphate was studied to determine Abstrart-SO, temperature dependencies of the reaction rate. The process was carried out at relatively high sulphuric acid (reaction product) concentrations with regard to its application for SO2 removal from waste gases. Variation of the apparent activation energy has been linked with alteration of reaction rate determining steps. INTRODUCTION
The catalytic influence of some transition metal ions on oxidation of SO, absorbed in aqueous solutions has been known since the 19th century. However, there arc at least two important reasons why research on this process is
still
carried
on.
One
is
its
complexity
solution into the reactor, mocouple for indicating mixture. Both the reactor tem were immersed in a within f O.l”C.
and a Pyrex jacketed thertemperature of the reacting and the gas conditioning syswater bath and thermostatted
involving
laborious studies on the process mechanism and another one is its practical importance for SO1 emission control. To develop such a technology sufficiently safe data are needed, particularly the optimum catalyst, its concentration, relative concentrations of reagents, reaction temperature etc. In this work we attempted to determine experimentally the activation energy of oxidation of SO, absorbed into the MnSO. aqueous solution hoping to gain additional information as to the reliability of the reaction model proposed prcviously[l]. Our intention was also to explain discrepancies in the scarce literature on the intlucncc of temperature upon the process. Some authors reported relatively high values of activation energy (Hoathcr et 01.[2]-27.3 kcallmol ? 4%, Huss[3]--19.8 + 0.7 kcallmol), whereas others indicated very low ones (Tarbutton et a/.[41 did not observe any measurable effect of temperature. the value evaluated from the data of Copson et al. [5] is lower than 2 kcal/mol). EXPERIMENTAL
Apparatus Experiments on SO2 oxidation were conducted in a semibatch foam reactor with constant flow of gaseous reactants, previously passed through a mixer and a humidifier. We have chosen this type of gas-liquid contact as it is directly transferable to industrial scale. A vertical Pyrex tube was used as the reactor (Fig. I), supplied at the bottom with the fine glass-frit as a gas distributor and, at the top with the PTFE lid with through-pipes. The pipes were coupled with a reflux corklenscr discharging after-reaction gases, tube for liquid sampling, separatory funnel applied for introducing initial or excessive (taken through the sampling tube) ‘Author to whom correspondence should be addressed.
tPart of this work was presented at the European Conference
of the Federation of European Chemical Societies: Chemical Pathways in the Environment, Palaiseau, France 1980.
ikiaten’als Sulphur dioxide. The supply of sulphur dioxide was from a technical cylinder placed outside through an intermediate steel bottle (2 I.) placed in the vicinity of the reactor and periodically loaded by distillation of liquid SO* from the outer cylinder. The gas was also cleaned by passing it through a silica gel column and then through a fine porous glass plate. Air. Molecular oxygen from air was used as an oxidizing agent. Air was sucked from the outside of the laboratory with a diaphragm pump and next passed through a silica gel column, cloth titer and fine porous plate. Manganous s&hate. The analytical grade reagent was from PPH POCH, Gliwicc. It was applied without any further purification. Water. All solutions were prepared with redistilled water. Procedure As soon as the bath temperature was fixed at the desired level the flow of air was turned on and then the reactor filled with the catalyst aqueous solution of known volume and Mn concentration. Simultaneously with the solution the proper stream of SO, was introduced into the flowing air and hence the oxidation was under way. From the start of supplying SOI a run was timed and successive liquid samples were withdrawn. Only 1 ml liquid portions were needed for analysis and the excess of the withdrawn liquid was returned to the reactor. The temperature of the reflux condenser was adjusted according to the programme worked out on the basis of preliminary experiments so as to keep the possibly constant volume of the reacting liquid, independently of the increasing H2S04 concentration with the progress of oxidation. Samples were analysed for H2S04 to get information on the extent of reaction and for Mn to allow correction for changes in the reacting solution volume. Alkalimetric
122
W.
PASIUK-BRONIKOWSKA and
A.
SOKOLOWSKI
Fig. 1. Schematic drawing of experimental apparatus: 1, pyrex tube 0.07m i.d.; 2, gas distributor/glass-friitl; 3, PTFE lid: 4, reflux condenser: 5, separatory funnel: 6, sampling tube; 7, thermocouple; 8, humidifier; 9, mixer of gaseous reagents; 10, rotameter; 11, Ueometer; 12, fine porous plate; 13, cloth filter; 14, silica gel column; 15, diaphragm pump. titration with metal masking and calorimetric determination in the presence of formaldoxime were employed respectively. Below fundamental parameters of experiments are specified: MuSO concentration SO? concentration flow of gaseous mixture liquid volume temperature
1.2 x lo-‘-0.18 mol/dm3 0.5-3 vol.% 0.1 l-O.44 dm’/s 0.134.25 dm’ 14.2-43.3’C.
RILWLTSANDDISCUSSION The rate of sulphuric acid production which is equal to the rate of SOz oxidation can be obtained from the equation r = (dnJdt)/
V = dc,/dt
- (dcJdt)(cJcM)
(1)
where n. = n.(t) and V= V(t) were not measured while c, = c,(t) and CM = c&t) were experimentally determined concentrations of sulphuric acid and manganese, respectively. The data points were expressed as higher order polynomial functions of time using the least squares method. Exemplary fitted curves are shown in Pi. 2. In view of previous reports[l-3.51 as well as of observations made in this work sulphuric acid causes significant retarding effect on the rate of its production. Therefore to examine the reaction sensitivity to temperature one should compare reaction rates at diRerent temperatures but alike acid concentrations or reaction extents when starting with an aqueous solution containing no acid. Temperature dependencies for SO2 oxidation at MnSOa concentration 5 x lo-” mol/dm3 and various
5
10
15 tno-T
Fii.
20
25
30
S
2. Sulpburic acid and manganese concentration 14.2’C; 0, 25.1”C and Cl, 34.8”C.
va time;
A,
reaction extents are plotted in Fig. 3. It shows results of experiments given as In r vs l/T instead of the typical Arrhenius plot, as the process is complex and the form of its rate equation may he still in question. Under the condition that experiments were performed at constant concentrations of both substrates the value of the apparent activation energy found from Fig. 3 did not differ from that determined from the classic Arrhenius equation. Data points reported in Fig. 3 are arranged in two regions. one of variable activation energy which
Activation energy variation for catalytic oxidation of aqueous SOI
I n
O,l
0.2 “_
-t
123
a3 .
Ob
I
dnP/S
<
Fi.
f
1
k
“’ L
I
32
, 3.3
3.L iO?T,
K-’
3.5
4. Dependence of the process rate upon gas flow rate for cM= 5 x IOFmolldm’and temperature 32.PC.
more probable. For this purpose experiments were carried out at all parameters constant but the gas Bow rate which was varied to change hydrodynamic conditions determining transfer of gaseous reactants. Figure 4 shows that at the gas flow rate 0.236 dm’ls applied in the main set of our experiments the influence of hydrodynamic conditions was not detectable. Rates of SO2 oxidation plotted in Fii. 4 are the maximum ones attained during runs at 32.7”C (found by extrapolation to c. = 0). These results allow more firmly to ascertain that the simultaneous absorption of oxygen and sulphur dioxide accompanied by the catalytic oxidation of the latter goes in the kinetic regime throughout tbe conditions described in Fii. 3 up to the temperature as in Pii. 4. However, further increase in the temperature might partly bring dilfusional limits on the examined oxidation process. The effect of temperature on the rate of SO2 oxidation at the relatively bigb MnSO., concentration (up to 0.18 mol/dm3) was rather weak. It is illustrated in Fig. 3 (dotted line) as derived from initial process rates of experiments with upper catalyst concentrations (see diffusional regime in Fig. 5). In this case the liquid side resistance of the oxygen transport was the rate -controlling phenomenon. Nevertheless, the foam reactor is not convenient for studying the process in the diffusional regime. It is mainly because the area of gas-liquid interface, especially foam height, is dependent upon the physical properties of the reacting mixture which change with the reaction going on. For SOa oxidation catalysed with MnSOd in aqueous solutions Pasiuk-Bronikowska and Bronikowski[l] have reported that the SO, absorption rate may be given by the empirical relation:
I
Pi. 3. Temperature dependencies for SO1 oxidation at c, = 5 x lo-’mol/dm’(exemplary points: + , c, = 0 mol/d&; A, c, = 0.408moI/dd; 0, E. = 0.816 mol/dm’). Values of activation energy at higher temperatures (left side of the diagram) are: 1.03. 2.75, 4.32. 6.10. 8.83. 10.70kcai/mol respectively with c. increasing.
decreases by several times with lowering sulphuric acid concentration and another one of almost constant activation energy 24 f 2.4 kcal/mol. The latter is extended to higher concentrations of sulpburic acid and lower temperatures. To get further information on the character of the regions it was necessary to answer whether they were diffusional or kinetic. Therefore fluxes of gaseous substrates (sulphur dioxide and oxygen) from the gas phase into the liquid bulk were calculated on the basis of previously found rates of oxidation as well as of the estimated mass transfer parameters k, = 5.5 x IO-’ m/s and (I = 14.5 x 10’rn-‘, and their values compared. The mass transfer parameters were determined from the correlations suggested by Gestrich et af.161. The comparison indicated that tbe slow reaction regime could be applied to both regions. The calculations also disclosed that with the additional increase in the rate of oxidation one might expect the appearance of the liquid side resistance first due to oxygen and not to sulphur dioxide. Oxygen diffusional limitation calculated on this basis is shown in Fii. 3 (dashed lime). As the estimation of kr and a was rather rough it seemed reasonable to verify practically the kinetic regime, especially for the variable activation energy region. In this instance the diffusional limitation wa*
llr= AdeM
+ B(c,/c,‘)
(2)
llr = A&,
+ B(c,/cM2)
(31
or by
depending on relative rates of appropriate step reactions. Here, the equations have been written in a converted form which allows to visualize that by examining a run
W. PASIUK-BRONIKOWSKAadd A. SOKOLOWSKI
124
where cm = [Mn”]([Mn3’] was negligible in comparison with the concentration of MnS04 throughout experiments), k, is a composite rate constant being a product of appropriate step reaction and equilibrium constants, and K’ and K” are equilibrium constants of particular step reactions. Comparison of eqn (3) with eqn (4) leads to A,= K’lk, and B = Wk,. When co= const which may be a rough approximation for a run eqn (4) becomes:
where k, = kc0 and K, = K”(co/cM2). As results from the equations contribution of the oxygen solubility temperature effect to the empirical value of the apparent activation energy should be considered. This effect is expected to be excluded when K,c, B K’. In this case lo-’
IO-'
1w2 cM
,
W, = (kJK”)cM’.
(6)
mol/dm’
Fii. 5. The influence of c,,, on the slope of the line l/r = f(c,) (eqns 2 or 3) disclosing diffusional limitation. as l/r vs c, the regime of gas absorption may be easily recognized. For both cases (eqns 2 and 3) when the kinetic regime (slow reaction) takes place the slope of a line should be inversely proportional to the squared M&O4 concentration (S = B/ca) provided all relevant parameters are kept constant. Similarly the intercept I = A Jc,’ (eqn 2) can be based on for the reaction going in the 0, independent region. The regularity vanishes in the diiusional regime. This is demonstrated in Fig. 5. It gives the slopes calculated for several runs performed at different values of cIK and temperatures 20.0”, 32.7’ and 40.0°C, plotted against c,,, (log-log coordinates). The data points follow the expected relationship up to a certain value of CM,being lower for the higher reaction temperature. Further increase of the MnS04 concentration indicates transition from the kinetic to the diiusional regime for which the above relationship ceases to hold. The results conlirm that conditions for all experiments shown in Fig. 3 below the dotted line corresponded to those of the kinetic regime. Even at the temperature 4O.O“Cthe intbtence of diffusion appeared only when cM >6~ lo-’ mol/dm” whereas experiments in Fii. 3 were at C~ = 5 x lo-“mol/dm3. Hence it can be concluded that the unusual effect of temperature observed in Fig. 3 reveals mainly the complex reaction mechanism. As the plot is curved downward this is not the case of parallel pathways but of reaction successive steps. According to Pasiuk-Bronikowska et a/.[11 who proposed the reaction model for oxidation of SOI catalysed with MnSOd the theoretical rate equation for the set of parameters applied in this work may be expressed as:
Plotting experimental data as rc, against c, the approached values of (t/K’9cM2 = const could be found for several temperatures in the lower range and hence the single apparent activation energy as given in Fig. 6 (a and b-l). Values of the coefficient (&,/K”)c,’ were also calculated from eqn (4) takine k&K’ as the initial rate of oxidation at c. = 0 (see b-2 in Fig. 7). Thus obtained values for the activation energy are 20.2 -(b-l) and 23.4 kcal/mol (b-2) indicating discrepancy caused by the way of data treatment. The results bear a slight compensating effect with respect to the second term of denominator in eqn (4) as co decreases with the rise of c.. Huss[31 who supplied the best evidence for his experiments gave the value of activation energy found at c0=7.S x lo-*mol/dm3, c, = 0 and c, = 3.31 x lo-’ mol/dm3 for temperatures between 25 and 38°C. In view of previous considerations on reaction mechanism[l] the conditions were fulfilled for the kinetics described by:
r = kocM2.
(7)
Therefore values of activation energy found in this work and reported by Huss should not necessarily be equal. CONCLUSIONS
Variation of activation energy for the reaction of SO? oxidation catalysed with Mn has been proved experimentally. To cbstinguish between the tirst order reaction and diiusion controlled kinetics with respect to oxygen the influence of other reagents (MnS04 and H2S04) in the complex reaction was successfully examined. Attempts were made to link such a behaviour with the reaction mechanism, which allowed to settle that SO* oxidation kinetics may or may not significantly depend on temperature according to the reaction order with respect to oxygen. With the order increasing (from 0 up to 1) the value of apparent activation energy diminishes from 19.8 kO.7 kcallmol in the absence of HJO, or 24%
Activation energy variation for catalytic oxidation of aqueous SO*
a
Fig. 6. Determination
24kcal/mol at relatively high to about 1 kcal/mol. The latter tors at high Mn concentrations overcome the energy-related acid contamination.
of apparent activation energy basing on eqn (6) (a and b-l) or eqn (4) (b-2).
concentrations of H2S04 is attained in foam reacencountered in practice to problem of environment
NOTATION
specific interfacial area with respect to gas-liquid volume, m-’ kinetic constants in experimental rate equations, mol s/dm’ and s respectively (eqns 3 and 4) kinetic constant in experimental rate equations (3 and 4), s manganous sulphate concentration, molldm3 oxygen concentration (in liquid), mol/dm3 sulphuric acid concentration, mol/dm’ apparent activation energy of the reaction (process), kcallmol intercept for I/r = f(c.) (oxygen independent region), dm3 s/m01 liquid side mass transfer
coefficient,
m/s
complex
rate constant
complex equilibrium constants moles of sulphuric acid, mol rate of the reaction, molldm’ s slope for i/r = f(c,), s/mol* dm6 time, s temperature of the reaction (process), “C absolute temperature, K volume of reacting liquid, dm3 gas flow-rate, m”ls
REFERENCF.S
Cl1 Pasiuk-Brouikowsks
Sci. 1981 36 215.
W. and Bronikowski T.. C/rem. Ensare - _
Dl Ho&her R. C. and Goodeve 30 11‘49.
C. F., Trans. Faraday
Sot. 1934
Huss A. Jr., Ph.D. Thesis. Urbana. Illinois 1978. E:; Tarbutton G., Driskell I. C., Jones T. hf., Gray F. J. and Smith C. M., bd. Engng Chem. 1957 49 392. PI Copson R. L. and Payne J. W., Ind. Engng Chem. 1933 25 9m ___. WI G&rich W., Esenwein H. and Kmuss W.. ht. Chem. Engng 1978 18 38.