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The oxidation rate of sulphide in sea water
Watt.r Re~ectrch. Voi. 8. pp. 395 to 41~). Pergamon Press 1974. Printed in Great Britain.
THE OXIDATION RATE OF SULPHIDE IN SEA WATER TOM ALMGRENand INGELAHAGSTROM Department of Analytical Chemistry, University of Gothenburg, FacL S-402 20 G6teborg 5. Sweden
(Received 19 October 1973) Abstract--Computer processed potentiometric measurements with a Ag/AgzS membrane electrode have been used to determine the order and rate of the oxidation of sulphide by oxygen in sea water at pH 8"0 and 8"5.The reaction is found to proceed through a pseudo first order mechanism when excess oxygen is present and the half-lives are calculated for different oxygen-sulphide ratios. The pH dependence is complex and probably influenced by catalysts. The microbial activity has been briefly studied and no changes in reaction rate were observed between sterilized and non-sterilized samples. Therefore bacteria do not probably take part in the rate-determining removal of sulphide.
INTRODUCTION
The occurrence of sulphide in the Baltic deep basins during the last decade and the generation of sulphide in heavy polluted recipients have caused a great deal of interest in the mechanisms of sulphide formation and removal in sea water. For the understanding of the chemistry of sulphide in anoxic waters it is important to know the rate constant for the reaction between oxygen and sulphide in sea water. Avrahami and Golding (1968) studied the reaction in alkaline aqueous solutions at pH 11-14. They found the reaction to be close to first order with respect to sulphide when other variables were constant. The rate-determining reaction step was found to be: H S - + {,O2---,SO2- + H + and the sulphite is rapidly removed to form thiosulphate and subsequently sulphate. They also studied the temperature dependence of the reaction and found a rate increase of 1.5 times for a 10°C increase in temperature. Also C)sdund and Alexalader (1963) found that the oxidation of small amounts of sulphide can be treated as a first order reaction in sea water saturated with oxygen. They reported a half-life for sulphide of about 20 min. Cline and Richards (1969) state that the oxidation in sea water has complex kinetics, but can be treated as a mixed second order reaction. The half-life of sulphide calculated from their Figure 2 is 15 h at 9-8°C. The half-life calculated from their rate constant is however 5,6 h for the same experiment. Sorokin (1971) studied the reaction in water from the active oxidation zone in the Black Sea. He found the reaction to be mainly chemical with the formation of sulphate and thiosulphate. Bacteria participate in a further oxidation of thiosulphate. The half-life of sulphide was according to his Figure 8 about 10 h. Chen and Morris (1970 and 1972a)investigated the initial velocities of w.a. 8,7--A
sulphide oxidation in fresh water and its pH dependence. They found an induction period for the reaction of 0.2-6 h. The reaction is suggested to proceed through a chain mechanism and the reaction order was found to be 1-34 with respect to sulphide. The pH dependence was very complex and the half-life for sulphide was about 50 h at pH 7-94. Chen and Morris (1972b) investigated the change in reaction rate due to catalyses and inhibition. They found the rate tO be very sensitive to catalysis by metal ions; especially Ni(ll) and Co(II). These previous results show considerable disagreement. Both the half-lives and the proposed mechanisms for the oxidation reveal great differences in experimental as well as in calculated results. Thus there is a need for further and perhaps simplified investigations in order to establish half-lives and kinetic models for use in sea water basins (e.g. in fjords and the Baltic .proper).
Choice of method A simple way to continuously monitor the sulphide concentration, is to measure the e.m.f, between a sulphide selective electrode and a calomel reference electrode. The method has the advantage of being very sensitive to the sulphide ion concentration, with a detection limit far below 10 -t5 M, provided the total sulphide concentration is not too low ( > 10 -6 M). Furthermore, the measurements can be carried out continuously in the reaction vessel after the addition of Na2S to aerated sea water. Pseudofirst order kinetics Assuming the rate-determining step: H S - + ~O_,--SO32- + H ÷ the reaction involves collisions between sulphide and 395
396
T. ALMGRENand 1. HAGSTROM
oxygen and the rate of sulphide disappearance can be described
dES(- II)] d~
=
k , , + , [ S ( - II)]"[O2]"
(I)
where k,,,~., denotes the rate-constant for the m + n order reaction, and S ( - I I ) is the sum S 2- + H S - + H_,S. At great excess of oxygen equation (1) can be approximated d [ S { - II)] = k , . [ S ( dt
II)]"
(2)
where k,. is a pseudo rate-constant. The usual way in determining reaction order and thus the rate constant is to fit the experimental data to an expression such as equation (2) with varied integer values of m. Assuming a pseudo first-order reaction, m = I, integrating equation (2) gives log[S(-II}] = log[S(-II)]o - k, .t/2.303
(3)
where [S(-II)]o is the initial total sulphide. If the reaction really corresponds to a pseudo first order treatment, a plot of log [ S ( - II)] vs time will produce a straight line if excess oxygen is present. From the slope -k,/2'303, k I can be determined. Assuming the, over-all reaction rate, m + n, to be 2 (bimolecular reaction), equation (l), gives k: k2 = kd[O2]. The half-life for a reactant is often a helpful constant, giving a more direct information about the rate and for a first order reaction
Keeping pH and temperature constant it is apparent that E. the e.m.f, of t.he cell. is directly proportional to log [ S ( - I I ) ] and thus a plot of E vs time produce the same slope as a first order plot. c.f. equation (3). EXPERIMENTAL
Apparatus To monitor the sulphide decrease a silver/silver sulphide membrane electrode, Orion Research 94-16, and a calomel reference, Radiometer K401, was used. A HP 2114B computer (cf. Anfalt and Jagner, 1971) was programmed to process the e.m.f, readings and to calculate the regression line through the e.m.f.-time values at the end of each run. The oxidation was carried out in a polyethylene beaker, closed by a Teflon lid. The reaction vessel was so constructed that it could be completely filled with sample and then closed to prevent air contamination. The sodium sulphide solution was introduced by an AGLA micrometer syringe, through a small hole in the Teflon lid which thereafter was closed by a rubber stopper.
Materials N a z S ' 9 H z O p.a., Merck No 6638, carefully rinsed with distilled water and dried between filter paper, was used for 0.1 M sulphide stock solution. Borate buffers, pH 8-0 and 8-5, were made from B(OH h p.a., Merck 165, and NaOH p.a., EKA. Standard Sea water, I.A.P.S.O. Charlottenlund, Denmark. Water from the Byfjord, a fjord at Uddevalla on the west coast of Sweden. Chloroform (1% ethanol) p.a., Fisher No C 374 was used to inhibit bacterial activity.
t½ = ln2/k,.
Analytical procedure
Initial velocities Looking at the very first part of the reaction
Ri = ( d[S(-II)])dt,= o = k , [ S ( - I I ) ] ,
(4)
is valid for different initial sulphide concentrations if the initial oxygen concentration is kept constant; k, being the initial rate-constant with respect to sulphide. The logarithmic form of equation (4) log Ri = log k, + m l o g [ S ( - II)].
(5)
shows that the reaction order m, is equal to the slope of a plot of log Ri vs log [ S ( - II)]o.
Electrode response A preliminary investigation showed that the sulphide electrode has a Nernstian response to the sulphide ion E = E ° + ½RTF - t In 101og[S2-].
The rate of sulphide decrease in samples of two different pH, 8.0 and 8-5, was measured for different initial sulphide concentrations. The e.m.f, between the indicator electrode and the reference was read each 5th min by the computer. The whole run was approximately 6 h, occasionally 10 h. The samples were prepared from standard sea water saturated with air, approx 200 teM at 2TC and 35%o salinity, borate buffer and sulphide solution. Precautions were taken to prevent air bubbles from remaining in the vessel. Magnetic stirring and computer readings were started as soon as the sulphide was introduced into the sample. The initial sulphide concentration was varied between 1 and 200 /~M. Several runs were made at same pH and initial sulphide concentration in order to check the precision of the measurement. To check the microbiological influence on the rate of reaction, water from the Byt~ord was used. Samples were taken from both oxygen-containing (above pycnocline) and anoxic
Oxidation of sulphide in sea-water (below pycnocline) waters. The anoxic samples were treated with great care to prevent oxygen contamination. The two samples were mixed in proportions yielding approx 30/~M sulphide. The rate of sulphide decrease was then compared with samples to which only buffer was added and samples to which both buffer and chloroform were added. The reaction was carried out at room temperature, and the temperature was checked now and then before and after a run. It was found to vary between 23 and 24~C, All equipment used for the experiments was carefully cleaned and washed in doubly distilled water. The silver/silver sulphide membrane electrode was polished approximately every 10th run with diamond paste, and thereafter cleaned with alcohol. This procedure was found to have a good effect on the electrode stability.
397
61C-
%*%%ae%
rain
Fig. 2. The e.m.f. IEmV) vs time (t rain) for 50 #.~t total initial sulphide concentration at pH 8.0.
300 Table 1. Reaction rates for the oxidation of sulphide in sea water saturated with oxygen pH
[S( -ll)]o #.,a
k I x 103 rain -I
t½min
8"0 8.0 8-0 8.0
l 5 6 10
18 8"5 7-5 6"0
8-0
20
5' 1
8.0 8-0 8.0 8'5 8.5 8.5 8'5 8.5
50 80 200 1 5 10 50 60
3.9 3-4 2.5 13
38 8l 92 120 140 180 200 280 53 63 74 120 130
II
9.3 5.8 5'2
2oc
~ Ioo
RESULTS
With the exception of the first 10-15 rain the e.m.f.time plots were in good agreement with the theory for a pseudo first order reaction, c.f. equation (3). Espe-
o
!
,oo
I
[o~o
,_
2oo
[s t-O)]o Fig. 3. The half-life (~ mini for different initial oxygen-sulphide ratios, at pH 8.0 (circles) and pH 8"5 (squares).
5~
18
54O
?
tG
o e x i
~20
tS 4 ql
2 G a
rain
Fig. 1. The e.m.f. (EmV) vs time (t min) for 1 ~M total initial sulphide concentration at pH 8.5.
I 150
I
Is t-m] o ~,. Fig. 4. The pseudo first-order constant (k~ min- t) for different initial sulphide concentrations, at p H 8-0 (circles) and pH 8-5 (squares).
398
T. AL~,IGRENand I. HAGSTROM was about + 3 min in t$. Figure 3 gives the half-life for the different oxygen sulphide ratios and Fig. 4 shows the variation of k, with initial sulphide concentration. The comparison between sterilized and non-sterilized samples showed no discernible difference in reaction rate. The half-life obtained, 74 rain. refers to approximately the ratio 20 between oxygen and sulphide. which is in a~eement with the half-life obtained in standard sea-water at pH 8.5. F i b r e s 5 and 6 are logarithm plots for initial velocities at pH 8-0 and 8.5 vs initial sulphide concentration. Both plots give a reaction order of about 1 with respect to sulphide. The difference between half-lives at pH 8-0 and 8.5 is plotted for each initial sulphide concentration in Fig. 7. It shows the complexity of the pH dependence of the reaction.
too [st-D~] o
Fig. 5. Initial velocity plot. pH 8"0.
Cg
DISCUSSION First order interpretation
,,q [si-m]. Fig. 6. Initial velocity plot, pH 8"5. cially the data from initial sulphide concentrations, 1t0 ltM, produced lines with correlation coefficients of about 0-99, c.f. Fig. l. At higher initial concentrations of sulphide the e.m.f.-time function tends to curve c.f. Fig. 2. Table I shows the results of a first order interpretation of the e.m.f.-time data. The maximum deviation from duplicate runs used for the interpretation
7O
:y
60
.~ --
30!
2°1- It o
Dependence on pH
• I
I
In the interpretation of the data as kt or t½ for different initial sulphide concentration it is assumed that the oxygen concentration is constant during the reaction. An initial sulphide concentrations greater than 20 ~ t there must be considerable decrease in the oxygen concentration. The deviation from straight-line relations for the e.m.f.-time plots at higher initial sulphide concentrations, c.f. Fig. 2, is most likely a result of this assumption. Nevertheless we preferred to interpret data in a single way for the sake of comparison. The half-lives obtained in this way are useful for calculations in the zone between oxic and anoxic water layers, where all kind of ratios between oxygen and sulphide can be expected. Figures 3 and 4 indicate that k t (and t½) will reach a constant value as the ratio of oxygen to sulphide increases. Therefore the kt value for I I~M sulphide (or ratio 200) can be regarded as a more or less true pseudo first order constant for the oxidation rate of sulphide in sea-water. From this constant an overall second order constant, k 2, can easily be calculated, by simply inserting the oxygen concentration, k_, = kt/(02) o, which yields 9-0 x 10 -5 #.~l rain - t at pH 8-0 and 6"5 × 10- 5 at 8"5.
I
I
I
5O
lit
Fig. 7. The difference in half-lives (At-~)between pH 8'0 and 8.5 at different initial sulphide concentrations.
Obviously the variation of rate with pH is dependent of the initial sulphide concentrations. A possible explanation for this is that some catalyst is present in low concentrations preferring lower pH. The rate increases normally with increasing pH due to increasing [ H S - ] / [ S ( I I ) ] ratio, but with a catalyst preferring low pH, this normal effect will be less pronounced at low initial sulphide concentration. Cben and Harris (1972)
Oxidation of sulphide in sea-water found very complex pH and concentration dependence also in fresh water.
Initial velocities The initial velocity interpretation is merely an additional check on the reaction order. In such an interpretation, there is no need for high oxygen excess, as only the I0 per cent of the reaction (except for the adjustment period) is used. for calculating the initial velocities. The results (c.f. Figs. 5 and 6) are a good indication on a first order reaction with respect to S ( III. m = 0-97 at pH 8"0 and 0.98 at pH 8.5. The anomaly at I pM initial sulphide concentration at pH 8"0 is probably due to the same catalytic factors as discussed under the pH dependence. Also at the two highest initial concentration an anomaly is apparent. In this case it is a loss in initial velocity. This effect could be caused by some inhibitor, e.g. produced as a reaction product.
The microbial influence As there is no doubt that bacteria take part in the oxidation reaction between sulphide and oxygen it is remarkable that we found no influence on the rate of reaction. The conclusion must be, that the bacteria shift the path after the initial rate-determining chemical reaction, thus producing the differences in the concentration of thiosulphate and sulphate observed by Sorokin (1971).
Sulphurformation It was occasionally observed that the solution became opalescent which was ascribed to the formation of elemental sulphur. Similar observations have
399
been made at the p.~cnoclines in B~fjorden qNyquisL 1972~ and in the Black Sea !Pilipchuk. 1973). The occurrence of sulphur in the reaction vessel does not alter the rate constant and sulphur is thus not a primary reaction product. Since both thiosulphate and sulphate are subsequently formed from sulphite the sulphur is most likely formed by spontaneous decomposition of thiosulphate
s_,o~- --, s o ~ - + s(s). In general this reaction is slow, but may probably be triggered by small impurities or bacteria.
Comparison with previous work Table 2 gives the results of other workers, in comparison with selected values from Table I. Three papers deal with the reaction in sea-water. C)stlund and Alexander (1963), Cline and Richards (1969) and Sorokin (1971). The ()stlund and Alexander values are very low, about seven times lower than comparable results by us. This can probably be assigned to the reaction conditions as they seem to ha,~e used an open vessel for the oxidation. The Sorokin value of about 10,000 rain, estimated by us from his data, in Black Sea pycnocline water seems very high. and can have been caused by some inhibition factor. Not knowing the initial oxygen concentration in his second experiment it is difficult to compare his t½ = 600 min calculated by us with the other works. It seems, however, to be more close to them than his first figure. The values of Cline and Richards are in the same order as our sulphide half-lives although the agreement could be better. Since the half-life calculated from the Cline and Richard rate constant differs very much. about three
Table 2. Comparison of sulphide oxidation rates in natural waters
[0., ]°/~M
[S( - 11)]o/~M
t(°C)
pH
r~min)
Constant saturated 200 240 120 60 12.5 +air 800 200 200 200 200
100 40 60 60 60 "- 6 14.7 100 50 200 50 200
-25 9-8 9.8 9.8 8-9 8-9 25 23--24 23-24 9-8 9'8
> 11 8.2 7.5-7.8 7-5-7.8 7.5-7.8 7.6-7.7 7.6-7.7 7.94 8"0 8-0 8"0 8.0
130
* Values corrected to 9.8"C.
24 175 336 861 10,000 600 3000 180 280 370* 570*
Order of reaction with respect to sulphide 1 1 1 1 1 ---
1.34 1 1 1 I
Reaction medium Fresh water Sea-water Sea-water Sea-water Sea-water Sea-water Sea-water Fresh water Sea-water Sea-water Sea-water Sea-water
Author Avrahami and Golding Ostlund and Alexander Cline and Richards Cline and Richards Cline and Richards Sorokin Sorokin Chen and Morris This work This work This work This work
400
T. ALMGRENand [. HAGSTRi)M
times lower, from a halt-life gained from a reaction plot in their work. it is. however, possible that all values from their work in Table 2 are three times too low. If so, the discrepancy between their values and ours is considerable. The t½-values of Chen and Morris points in that direction. They used a method similar to Cline's 11964) to determine the sulphide concentrations. They obtained a t½-value of 3000 min, which is very high in comparison with ours. Although their experiments were made in fresh water solution, that is without all the possible catalysts in sea-water, this fact might indicate that there is an analytical difference between this photometric method and electrometric analysis of sulphide. The value of Avrahami and Golding (1968) is also hard to compare due to the high pH. They could use a direct photometric determination of the concentration of H S - . The reaction order reported is close to first order with respect to sulphide. Cline and Richards (1969) and Chen and Morris (1972)report very complex kinetics with no single rate equations. Avrahami and Golding (1968), C)stlund and Alexander (1963) and we have found more simple kinetics and more close to a pseudo first order reaction with respect to sulphide.
Acknowledgements--The authors wish to thank professor David Dryssen for valuable discussions. Part of this work was carried out when one of the authors (T.A.I was
employed by the Fishery Board of Sweden. The work on Byljorden at our department is supported by the Swedish Environmental Protection Board and the work on chemistry of sea water by the Swedish Natural Science Research Council. The computer was purchased by a grant from the foundation of Knut and Alice Wallenberg. REFERENCES
Anf'zilt T. and Jagner D. [1971) A computer processed semiautomatic titrator for high precision analysis. Anal. Chim. Acta 57, 177-183. Avrahami M. and Golding R. M. (1968} The oxidation of the sulphide ion at very low concentrations in aqueous solutions. J. Chem. Soc. (A) 647-651. Chen K. Y. and Morris J. C. (1970) Oxidation of aqueous sulfide by O.,. 5th Int. Conf. War. Pollut. Res.. San Francisco Cali]~ Chen K. Y. and Morris J. C. (1972) Kinetics of oxidation of aqueous sulfide by O,. Enciron. Sci. Technol. 6, 529-537. Cline J. D. (1969) Spectrophotometric determination of hydrogen sulfide in natural waters. Limnol. and Oceanogr. 14, 454-458. Cline J. D. and Richards F. A. (1969) Oxygenation of hydrogen sulfide in sea water at constant salinity, temperature and pH. Environ. Sci. Technol. 3, 838-843. Nyquist G. (1972) personal communication. Pilipchuk M. F. (1973) personal communication. Sorokin Yu. I. (1971] Experimental data on the oxidation rate of hydrogen sulphide in the Black Sea. Okeanologija. ! I, 423-431. Ostlund G. H. and Alexander J. (1963) Oxidation rate of sulfide in sea water. A preliminary study. J. Geophys. Res. 68, 3995-3997.
THE OXIDATION RATE OF SULPHIDE IN SEA WATER TOM ALMGRENand INGELAHAGSTROM Department of Analytical Chemistry, University of Gothenburg, FacL S-402 20 G6teborg 5. Sweden
(Received 19 October 1973) Abstract--Computer processed potentiometric measurements with a Ag/AgzS membrane electrode have been used to determine the order and rate of the oxidation of sulphide by oxygen in sea water at pH 8"0 and 8"5.The reaction is found to proceed through a pseudo first order mechanism when excess oxygen is present and the half-lives are calculated for different oxygen-sulphide ratios. The pH dependence is complex and probably influenced by catalysts. The microbial activity has been briefly studied and no changes in reaction rate were observed between sterilized and non-sterilized samples. Therefore bacteria do not probably take part in the rate-determining removal of sulphide.
INTRODUCTION
The occurrence of sulphide in the Baltic deep basins during the last decade and the generation of sulphide in heavy polluted recipients have caused a great deal of interest in the mechanisms of sulphide formation and removal in sea water. For the understanding of the chemistry of sulphide in anoxic waters it is important to know the rate constant for the reaction between oxygen and sulphide in sea water. Avrahami and Golding (1968) studied the reaction in alkaline aqueous solutions at pH 11-14. They found the reaction to be close to first order with respect to sulphide when other variables were constant. The rate-determining reaction step was found to be: H S - + {,O2---,SO2- + H + and the sulphite is rapidly removed to form thiosulphate and subsequently sulphate. They also studied the temperature dependence of the reaction and found a rate increase of 1.5 times for a 10°C increase in temperature. Also C)sdund and Alexalader (1963) found that the oxidation of small amounts of sulphide can be treated as a first order reaction in sea water saturated with oxygen. They reported a half-life for sulphide of about 20 min. Cline and Richards (1969) state that the oxidation in sea water has complex kinetics, but can be treated as a mixed second order reaction. The half-life of sulphide calculated from their Figure 2 is 15 h at 9-8°C. The half-life calculated from their rate constant is however 5,6 h for the same experiment. Sorokin (1971) studied the reaction in water from the active oxidation zone in the Black Sea. He found the reaction to be mainly chemical with the formation of sulphate and thiosulphate. Bacteria participate in a further oxidation of thiosulphate. The half-life of sulphide was according to his Figure 8 about 10 h. Chen and Morris (1970 and 1972a)investigated the initial velocities of w.a. 8,7--A
sulphide oxidation in fresh water and its pH dependence. They found an induction period for the reaction of 0.2-6 h. The reaction is suggested to proceed through a chain mechanism and the reaction order was found to be 1-34 with respect to sulphide. The pH dependence was very complex and the half-life for sulphide was about 50 h at pH 7-94. Chen and Morris (1972b) investigated the change in reaction rate due to catalyses and inhibition. They found the rate tO be very sensitive to catalysis by metal ions; especially Ni(ll) and Co(II). These previous results show considerable disagreement. Both the half-lives and the proposed mechanisms for the oxidation reveal great differences in experimental as well as in calculated results. Thus there is a need for further and perhaps simplified investigations in order to establish half-lives and kinetic models for use in sea water basins (e.g. in fjords and the Baltic .proper).
Choice of method A simple way to continuously monitor the sulphide concentration, is to measure the e.m.f, between a sulphide selective electrode and a calomel reference electrode. The method has the advantage of being very sensitive to the sulphide ion concentration, with a detection limit far below 10 -t5 M, provided the total sulphide concentration is not too low ( > 10 -6 M). Furthermore, the measurements can be carried out continuously in the reaction vessel after the addition of Na2S to aerated sea water. Pseudofirst order kinetics Assuming the rate-determining step: H S - + ~O_,--SO32- + H ÷ the reaction involves collisions between sulphide and 395
396
T. ALMGRENand 1. HAGSTROM
oxygen and the rate of sulphide disappearance can be described
dES(- II)] d~
=
k , , + , [ S ( - II)]"[O2]"
(I)
where k,,,~., denotes the rate-constant for the m + n order reaction, and S ( - I I ) is the sum S 2- + H S - + H_,S. At great excess of oxygen equation (1) can be approximated d [ S { - II)] = k , . [ S ( dt
II)]"
(2)
where k,. is a pseudo rate-constant. The usual way in determining reaction order and thus the rate constant is to fit the experimental data to an expression such as equation (2) with varied integer values of m. Assuming a pseudo first-order reaction, m = I, integrating equation (2) gives log[S(-II}] = log[S(-II)]o - k, .t/2.303
(3)
where [S(-II)]o is the initial total sulphide. If the reaction really corresponds to a pseudo first order treatment, a plot of log [ S ( - II)] vs time will produce a straight line if excess oxygen is present. From the slope -k,/2'303, k I can be determined. Assuming the, over-all reaction rate, m + n, to be 2 (bimolecular reaction), equation (l), gives k: k2 = kd[O2]. The half-life for a reactant is often a helpful constant, giving a more direct information about the rate and for a first order reaction
Keeping pH and temperature constant it is apparent that E. the e.m.f, of t.he cell. is directly proportional to log [ S ( - I I ) ] and thus a plot of E vs time produce the same slope as a first order plot. c.f. equation (3). EXPERIMENTAL
Apparatus To monitor the sulphide decrease a silver/silver sulphide membrane electrode, Orion Research 94-16, and a calomel reference, Radiometer K401, was used. A HP 2114B computer (cf. Anfalt and Jagner, 1971) was programmed to process the e.m.f, readings and to calculate the regression line through the e.m.f.-time values at the end of each run. The oxidation was carried out in a polyethylene beaker, closed by a Teflon lid. The reaction vessel was so constructed that it could be completely filled with sample and then closed to prevent air contamination. The sodium sulphide solution was introduced by an AGLA micrometer syringe, through a small hole in the Teflon lid which thereafter was closed by a rubber stopper.
Materials N a z S ' 9 H z O p.a., Merck No 6638, carefully rinsed with distilled water and dried between filter paper, was used for 0.1 M sulphide stock solution. Borate buffers, pH 8-0 and 8-5, were made from B(OH h p.a., Merck 165, and NaOH p.a., EKA. Standard Sea water, I.A.P.S.O. Charlottenlund, Denmark. Water from the Byfjord, a fjord at Uddevalla on the west coast of Sweden. Chloroform (1% ethanol) p.a., Fisher No C 374 was used to inhibit bacterial activity.
t½ = ln2/k,.
Analytical procedure
Initial velocities Looking at the very first part of the reaction
Ri = ( d[S(-II)])dt,= o = k , [ S ( - I I ) ] ,
(4)
is valid for different initial sulphide concentrations if the initial oxygen concentration is kept constant; k, being the initial rate-constant with respect to sulphide. The logarithmic form of equation (4) log Ri = log k, + m l o g [ S ( - II)].
(5)
shows that the reaction order m, is equal to the slope of a plot of log Ri vs log [ S ( - II)]o.
Electrode response A preliminary investigation showed that the sulphide electrode has a Nernstian response to the sulphide ion E = E ° + ½RTF - t In 101og[S2-].
The rate of sulphide decrease in samples of two different pH, 8.0 and 8-5, was measured for different initial sulphide concentrations. The e.m.f, between the indicator electrode and the reference was read each 5th min by the computer. The whole run was approximately 6 h, occasionally 10 h. The samples were prepared from standard sea water saturated with air, approx 200 teM at 2TC and 35%o salinity, borate buffer and sulphide solution. Precautions were taken to prevent air bubbles from remaining in the vessel. Magnetic stirring and computer readings were started as soon as the sulphide was introduced into the sample. The initial sulphide concentration was varied between 1 and 200 /~M. Several runs were made at same pH and initial sulphide concentration in order to check the precision of the measurement. To check the microbiological influence on the rate of reaction, water from the Byt~ord was used. Samples were taken from both oxygen-containing (above pycnocline) and anoxic
Oxidation of sulphide in sea-water (below pycnocline) waters. The anoxic samples were treated with great care to prevent oxygen contamination. The two samples were mixed in proportions yielding approx 30/~M sulphide. The rate of sulphide decrease was then compared with samples to which only buffer was added and samples to which both buffer and chloroform were added. The reaction was carried out at room temperature, and the temperature was checked now and then before and after a run. It was found to vary between 23 and 24~C, All equipment used for the experiments was carefully cleaned and washed in doubly distilled water. The silver/silver sulphide membrane electrode was polished approximately every 10th run with diamond paste, and thereafter cleaned with alcohol. This procedure was found to have a good effect on the electrode stability.
397
61C-
%*%%ae%
rain
Fig. 2. The e.m.f. IEmV) vs time (t rain) for 50 #.~t total initial sulphide concentration at pH 8.0.
300 Table 1. Reaction rates for the oxidation of sulphide in sea water saturated with oxygen pH
[S( -ll)]o #.,a
k I x 103 rain -I
t½min
8"0 8.0 8-0 8.0
l 5 6 10
18 8"5 7-5 6"0
8-0
20
5' 1
8.0 8-0 8.0 8'5 8.5 8.5 8'5 8.5
50 80 200 1 5 10 50 60
3.9 3-4 2.5 13
38 8l 92 120 140 180 200 280 53 63 74 120 130
II
9.3 5.8 5'2
2oc
~ Ioo
RESULTS
With the exception of the first 10-15 rain the e.m.f.time plots were in good agreement with the theory for a pseudo first order reaction, c.f. equation (3). Espe-
o
!
,oo
I
[o~o
,_
2oo
[s t-O)]o Fig. 3. The half-life (~ mini for different initial oxygen-sulphide ratios, at pH 8.0 (circles) and pH 8"5 (squares).
5~
18
54O
?
tG
o e x i
~20
tS 4 ql
2 G a
rain
Fig. 1. The e.m.f. (EmV) vs time (t min) for 1 ~M total initial sulphide concentration at pH 8.5.
I 150
I
Is t-m] o ~,. Fig. 4. The pseudo first-order constant (k~ min- t) for different initial sulphide concentrations, at p H 8-0 (circles) and pH 8-5 (squares).
398
T. AL~,IGRENand I. HAGSTROM was about + 3 min in t$. Figure 3 gives the half-life for the different oxygen sulphide ratios and Fig. 4 shows the variation of k, with initial sulphide concentration. The comparison between sterilized and non-sterilized samples showed no discernible difference in reaction rate. The half-life obtained, 74 rain. refers to approximately the ratio 20 between oxygen and sulphide. which is in a~eement with the half-life obtained in standard sea-water at pH 8.5. F i b r e s 5 and 6 are logarithm plots for initial velocities at pH 8-0 and 8.5 vs initial sulphide concentration. Both plots give a reaction order of about 1 with respect to sulphide. The difference between half-lives at pH 8-0 and 8.5 is plotted for each initial sulphide concentration in Fig. 7. It shows the complexity of the pH dependence of the reaction.
too [st-D~] o
Fig. 5. Initial velocity plot. pH 8"0.
Cg
DISCUSSION First order interpretation
,,q [si-m]. Fig. 6. Initial velocity plot, pH 8"5. cially the data from initial sulphide concentrations, 1t0 ltM, produced lines with correlation coefficients of about 0-99, c.f. Fig. l. At higher initial concentrations of sulphide the e.m.f.-time function tends to curve c.f. Fig. 2. Table I shows the results of a first order interpretation of the e.m.f.-time data. The maximum deviation from duplicate runs used for the interpretation
7O
:y
60
.~ --
30!
2°1- It o
Dependence on pH
• I
I
In the interpretation of the data as kt or t½ for different initial sulphide concentration it is assumed that the oxygen concentration is constant during the reaction. An initial sulphide concentrations greater than 20 ~ t there must be considerable decrease in the oxygen concentration. The deviation from straight-line relations for the e.m.f.-time plots at higher initial sulphide concentrations, c.f. Fig. 2, is most likely a result of this assumption. Nevertheless we preferred to interpret data in a single way for the sake of comparison. The half-lives obtained in this way are useful for calculations in the zone between oxic and anoxic water layers, where all kind of ratios between oxygen and sulphide can be expected. Figures 3 and 4 indicate that k t (and t½) will reach a constant value as the ratio of oxygen to sulphide increases. Therefore the kt value for I I~M sulphide (or ratio 200) can be regarded as a more or less true pseudo first order constant for the oxidation rate of sulphide in sea-water. From this constant an overall second order constant, k 2, can easily be calculated, by simply inserting the oxygen concentration, k_, = kt/(02) o, which yields 9-0 x 10 -5 #.~l rain - t at pH 8-0 and 6"5 × 10- 5 at 8"5.
I
I
I
5O
lit
Fig. 7. The difference in half-lives (At-~)between pH 8'0 and 8.5 at different initial sulphide concentrations.
Obviously the variation of rate with pH is dependent of the initial sulphide concentrations. A possible explanation for this is that some catalyst is present in low concentrations preferring lower pH. The rate increases normally with increasing pH due to increasing [ H S - ] / [ S ( I I ) ] ratio, but with a catalyst preferring low pH, this normal effect will be less pronounced at low initial sulphide concentration. Cben and Harris (1972)
Oxidation of sulphide in sea-water found very complex pH and concentration dependence also in fresh water.
Initial velocities The initial velocity interpretation is merely an additional check on the reaction order. In such an interpretation, there is no need for high oxygen excess, as only the I0 per cent of the reaction (except for the adjustment period) is used. for calculating the initial velocities. The results (c.f. Figs. 5 and 6) are a good indication on a first order reaction with respect to S ( III. m = 0-97 at pH 8"0 and 0.98 at pH 8.5. The anomaly at I pM initial sulphide concentration at pH 8"0 is probably due to the same catalytic factors as discussed under the pH dependence. Also at the two highest initial concentration an anomaly is apparent. In this case it is a loss in initial velocity. This effect could be caused by some inhibitor, e.g. produced as a reaction product.
The microbial influence As there is no doubt that bacteria take part in the oxidation reaction between sulphide and oxygen it is remarkable that we found no influence on the rate of reaction. The conclusion must be, that the bacteria shift the path after the initial rate-determining chemical reaction, thus producing the differences in the concentration of thiosulphate and sulphate observed by Sorokin (1971).
Sulphurformation It was occasionally observed that the solution became opalescent which was ascribed to the formation of elemental sulphur. Similar observations have
399
been made at the p.~cnoclines in B~fjorden qNyquisL 1972~ and in the Black Sea !Pilipchuk. 1973). The occurrence of sulphur in the reaction vessel does not alter the rate constant and sulphur is thus not a primary reaction product. Since both thiosulphate and sulphate are subsequently formed from sulphite the sulphur is most likely formed by spontaneous decomposition of thiosulphate
s_,o~- --, s o ~ - + s(s). In general this reaction is slow, but may probably be triggered by small impurities or bacteria.
Comparison with previous work Table 2 gives the results of other workers, in comparison with selected values from Table I. Three papers deal with the reaction in sea-water. C)stlund and Alexander (1963), Cline and Richards (1969) and Sorokin (1971). The ()stlund and Alexander values are very low, about seven times lower than comparable results by us. This can probably be assigned to the reaction conditions as they seem to ha,~e used an open vessel for the oxidation. The Sorokin value of about 10,000 rain, estimated by us from his data, in Black Sea pycnocline water seems very high. and can have been caused by some inhibition factor. Not knowing the initial oxygen concentration in his second experiment it is difficult to compare his t½ = 600 min calculated by us with the other works. It seems, however, to be more close to them than his first figure. The values of Cline and Richards are in the same order as our sulphide half-lives although the agreement could be better. Since the half-life calculated from the Cline and Richard rate constant differs very much. about three
Table 2. Comparison of sulphide oxidation rates in natural waters
[0., ]°/~M
[S( - 11)]o/~M
t(°C)
pH
r~min)
Constant saturated 200 240 120 60 12.5 +air 800 200 200 200 200
100 40 60 60 60 "- 6 14.7 100 50 200 50 200
-25 9-8 9.8 9.8 8-9 8-9 25 23--24 23-24 9-8 9'8
> 11 8.2 7.5-7.8 7-5-7.8 7.5-7.8 7.6-7.7 7.6-7.7 7.94 8"0 8-0 8"0 8.0
130
* Values corrected to 9.8"C.
24 175 336 861 10,000 600 3000 180 280 370* 570*
Order of reaction with respect to sulphide 1 1 1 1 1 ---
1.34 1 1 1 I
Reaction medium Fresh water Sea-water Sea-water Sea-water Sea-water Sea-water Sea-water Fresh water Sea-water Sea-water Sea-water Sea-water
Author Avrahami and Golding Ostlund and Alexander Cline and Richards Cline and Richards Cline and Richards Sorokin Sorokin Chen and Morris This work This work This work This work
400
T. ALMGRENand [. HAGSTRi)M
times lower, from a halt-life gained from a reaction plot in their work. it is. however, possible that all values from their work in Table 2 are three times too low. If so, the discrepancy between their values and ours is considerable. The t½-values of Chen and Morris points in that direction. They used a method similar to Cline's 11964) to determine the sulphide concentrations. They obtained a t½-value of 3000 min, which is very high in comparison with ours. Although their experiments were made in fresh water solution, that is without all the possible catalysts in sea-water, this fact might indicate that there is an analytical difference between this photometric method and electrometric analysis of sulphide. The value of Avrahami and Golding (1968) is also hard to compare due to the high pH. They could use a direct photometric determination of the concentration of H S - . The reaction order reported is close to first order with respect to sulphide. Cline and Richards (1969) and Chen and Morris (1972)report very complex kinetics with no single rate equations. Avrahami and Golding (1968), C)stlund and Alexander (1963) and we have found more simple kinetics and more close to a pseudo first order reaction with respect to sulphide.
Acknowledgements--The authors wish to thank professor David Dryssen for valuable discussions. Part of this work was carried out when one of the authors (T.A.I was
employed by the Fishery Board of Sweden. The work on Byljorden at our department is supported by the Swedish Environmental Protection Board and the work on chemistry of sea water by the Swedish Natural Science Research Council. The computer was purchased by a grant from the foundation of Knut and Alice Wallenberg. REFERENCES
Anf'zilt T. and Jagner D. [1971) A computer processed semiautomatic titrator for high precision analysis. Anal. Chim. Acta 57, 177-183. Avrahami M. and Golding R. M. (1968} The oxidation of the sulphide ion at very low concentrations in aqueous solutions. J. Chem. Soc. (A) 647-651. Chen K. Y. and Morris J. C. (1970) Oxidation of aqueous sulfide by O.,. 5th Int. Conf. War. Pollut. Res.. San Francisco Cali]~ Chen K. Y. and Morris J. C. (1972) Kinetics of oxidation of aqueous sulfide by O,. Enciron. Sci. Technol. 6, 529-537. Cline J. D. (1969) Spectrophotometric determination of hydrogen sulfide in natural waters. Limnol. and Oceanogr. 14, 454-458. Cline J. D. and Richards F. A. (1969) Oxygenation of hydrogen sulfide in sea water at constant salinity, temperature and pH. Environ. Sci. Technol. 3, 838-843. Nyquist G. (1972) personal communication. Pilipchuk M. F. (1973) personal communication. Sorokin Yu. I. (1971] Experimental data on the oxidation rate of hydrogen sulphide in the Black Sea. Okeanologija. ! I, 423-431. Ostlund G. H. and Alexander J. (1963) Oxidation rate of sulfide in sea water. A preliminary study. J. Geophys. Res. 68, 3995-3997.