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RESEARCH NOTES Effects of Platinum Electrode on Hydrogen, Oxygen, and Hydrogen Peroxide Formation in Aqueous Phase Pulsed Corona Electrical Discharge Michael J. Kirkpatrick and Bruce R. Locke* Department of Chemical and Biomedical Engineering, FAMU-FSU College of Engineering, Florida State UniVersity, Tallahassee, Florida 32310
Platinum, when used as the high-voltage needle electrode in pulsed corona electrical discharge in water, reduces the production rates of molecular hydrogen, hydrogen peroxide, and molecular oxygen by the discharge, in comparison to the production rates of those species when a nickel-chromium electrode is used. Reactions between hydrogen and oxygen and hydrogen and hydrogen peroxide are proposed to explain this result. It is proposed that these reactions occur on the surfaces of particles sputtered from the discharge electrode. 1. Introduction A pulsed corona electrical discharge (PCED) in water in the range of 1 J/pulse can have significant physical and chemical effects on the medium in which it propagates, including the production of stable active chemical species (H2O2, H2, O2),1,2 hydroxyl radicals,3,4 and O atoms,5,6 and the formation of shock waves.7 PCED may promote interactions between the active species created by the discharge with solid particles suspended in solution such as activated carbon.8 It has also been observed that, under some conditions, the materials used for electrodes in PCED in water can have effects on the degradation of organic compounds such as polychlorinated biphenyls.9 This suggests the possibility that reactions may occur on surfaces associated with the high-voltage electrode. The overall stoichiometry of active species formed by PCED in water follows that of reaction 1,1 and this stoichiometry is independent of reactor conditions such as discharge power or solution conductivity.
6H2O f 4H2 + 2H2O2 + O2
(1)
The rates of production of hydrogen and hydrogen peroxide were observed to be dependent on discharge power and solution conductivity, which is a result that was previously well-known for the case of hydrogen peroxide.2 The present work involves the effect of the high-voltage-electrode material on the chemical yields of hydrogen, hydrogen peroxide, and oxygen. 2. Experimental Methods The power supply and pulse-forming network used to produce the pulsed corona discharge in water utilized a rotating sparkgap-type apparatus and has been fully described in previous publications.1,3,8 The voltage rise time is on the order of ∼20 ns, and the current pulse width is on the order of ∼1 µs. The * To whom correspondence should be addressed. Tel.: (850) 410 6165. E-mail:
[email protected].
pulse width is dependent on solution conductivity; however, in this work, only one conductivity value was utilized (150 µS/ cm). Molecular hydrogen and oxygen contents were measured by gas chromatography (GC), as previously reported.1 The measurement of the hydrogen concentration was accomplished using nitrogen as both a reactor purge gas and a GC carrier gas, whereas helium was used for the purge and GC carrier gas for measurement of the oxygen concentration. Different purge/ carrier gases were used for the measurement of hydrogen and oxygen content, to maximize the sensitivity of the analysis of the two gases. The reason for using the same gas for the purge gas as that which was used for the GC carrier gas was simply to avoid large solvent peaks on the chromatograms. The different purge/carrier gases that were used for the hydrogen and oxygen measurements was the sole difference between the two sets of experiments. The hydrogen peroxide content was measured via absorption spectroscopy of a titania-hydrogen peroxide complex at 410 nm;10 this measurement was checked by titration with standardized potassium permanganate.11 The electrode configuration was point-to-plane, and the reactor consisted of a 1-L cylindrical Pyrex vessel; except for the use of a purge gas, the experimental arrangement was very similar to what which has been previously reported in more detail.3,8 In all the experiments presented here, positive-polarity high-voltage pulses were applied to the point electrode. Although no experiments were conducted using negative polarity, a different outcome with regard to hydrogen and oxygen production may be likely, in light of the fact that hydrogen peroxide production has been observed to be lower with negative polarity than with positive polarity.2 Discharge energy and pulse repetition frequency were the same for all of the experiments (∼1.1 J/pulse and 60 Hz, respectively), giving a power of 66 W. 3. Results and Discussion 3.1. Effect of Platinum Electrode on Hydrogen and Hydrogen Peroxide Production Rates. Figure 1 shows the average (from at least three experiments) hydrogen and hydrogen peroxide concentrations resulting from PCED in water for
10.1021/ie0511480 CCC: $33.50 © 2006 American Chemical Society Published on Web 02/10/2006
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Figure 1. Hydrogen and hydrogen peroxide production by electrical discharge in water with two types of needle discharge electrodes. The symbols correspond as follows: with the NiCr electrode, ([) H2 and (9) H2O2; and with the platinum electrode, (]) H2, and (0) H2O2.
Figure 2. ([) Hydrogen and (9) hydrogen peroxide production rates calculated from the data shown in Figure 1 for the nickel-chromium electrode.
nickel-chromium alloy and platinum high-voltage electrodes. The data for the nickel-chromium electrode is the same as that which appears in ref 1, and that information is presented here only for comparison; the new results presented in this communication concern only the use of platinum as a high-voltage needle electrode. The discharge voltage, frequency, and power, as well as the solution conductivities and temperatures, were the same for both cases. The only difference between the two cases is the material used for the high-voltage needle electrodes. As shown in the figure, for nickel-chromium, the hydrogen and hydrogen peroxide production rates were stable (for hydrogen, an initial increase is observed, because of mixing effects) over the course of 60 min. The rate of hydrogen peroxide production is the slope of the line shown in Figure 1, converted to units of mol/s, and the rates of hydrogen and oxygen production are obtained from the concentrations shown in Figure 1, multiplied by the total gas flow rate and converted to units of mol/s. Figure 2 shows the production rates calculated from the data shown in Figure 1, and it can be seen that when nickelchromium is used as a high-voltage needle electrode, the production rates of both hydrogen and hydrogen peroxide were stable and maintained a molar ratio of hydrogen to hydrogen peroxide of ∼2:1. However, when platinum is used as the highvoltage electrode, the production rates and the ratios of production rates change significantly, as shown in Figure 3.
Figure 3. ([) Hydrogen and (9) hydrogen peroxide rates calculated from the data shown in Figure 1 for the platinum electrode.
During the first 10 minutes of the experiment, the production rates of both species are just less than those observed for the nickel-chromium electrode; however, by the end of the experiment, both rates have dropped well below those observed for the nickel-chromium electrode. Specifically, the rate of hydrogen peroxide production decreases to approximately half of that observed with nickel-chromium, and the hydrogen production decreases to slightly more than one-third of that observed with the nickel-chromium electrode. The fact that the initial rates are similar for the two cases indicates that the “true” production rates are the same for both cases, and, therefore, the primary effect of platinum is to enhance the degradation of hydrogen and hydrogen peroxide, leading to smaller measured or “net” rates of production. Because the number of reactive species in this system is limited to compounds that contain hydrogen and oxygen, the number of chemical pathways that could possibly explain this degradation is limited. Considering only hydrogen and hydrogen peroxide, the overall reaction pathway that is depicted by reaction 2 may be postulated: Pt
H2 + H2O2 98 2H2O
(2)
However, because the changes in the molar production rates that are caused by the platinum electrode are not the same for hydrogen and hydrogen peroxide, reaction 2 alone cannot explain the data.
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Figure 4. Oxygen and hydrogen peroxide production with NiCr and Pt electrodes. The symbols correspond as follows: with the NiCr electrode, ([) O2 and (9) H2O2; and with the platinum electrode, (]) O2 and (0) H2O2.
3.2. Effect of Platinum on the Production Rates of Oxygen and Hydrogen Peroxide. As shown in Figure 4, the production of molecular oxygen is reduced by the use of platinum as the high-voltage needle electrode. Although the data for oxygen measurement are quite noisy (for reasons that have been discussed in more detail in ref 1), the basic result of a decrease in oxygen production is evident. The first two columns of Table 1 list the rates of production of the three species calculated from the average of the concentration data from the last 20 min of the experiments. There are two rows showing the rate of hydrogen peroxide production, because it was measured in both the experiments for hydrogen measurement and oxygen measurement. A difference in the production rate of hydrogen peroxide appears between the two cases; there is currently no explanation for this difference. For the case of the nickelchromium electrode (the first column), the rates appear roughly in the ratio suggested by reaction 1 (the rate of hydrogen production is twice that of hydrogen peroxide, which, in turn, is approximately twice that of the rate of oxygen production), with the caveat that the number for oxygen is affected by a large measurement error. The second column in Table 1 shows the observed rates of production with platinum as the point electrode, and the third column shows the difference between the first two columns, representing the effect of the platinum electrode. The net decrease in hydrogen peroxide production is on the order of half the net decrease in hydrogen production, indicating that reaction 2 can account for only about half of the effect of the platinum electrode on hydrogen. Therefore, there must be another explanation for why the hydrogen production rate decreases so much. A second overall reaction pathway (reaction 3) may be postulated to account for the remaining change in hydrogen production: Pt
2H2 + O2 98 2H2O
(3)
Considering the changes in hydrogen and hydrogen peroxide production from the hydrogen measurement experiments (the first two rows of Table 1) and the change in oxygen production shown in the third row of Table 1, only 4.4 µmol/s out of 7.3 µmol/s (the addition of the 3.0 µmol/s of reduction in the molar rate of hydrogen peroxide plus twice the 0.7 µmol/s of reduction in the rate of oxygen production) of the observed rate reduction in hydrogen production may be explained (this calculation assumes that oxygen production in the experiment shown in
Table 1. High-Voltage-Electrode Effect on Overall Production Rates of Hydrogen, Hydrogen Peroxide, and Oxygen Observed during the Last 20 Minutes of Experiments Production Rate (µmol/s) compound
nickel-chromium
platinum
net change
hydrogen hydrogen peroxide (hydrogen exp.) oxygen hydrogen peroxide (oxygen exp.)
11.8 6.2
4.5 3.2
-7.3 -3.0
2.3 6.0
1.6 0.6
-0.7 -5.4
Figure 1 changed in the same way as that observed for the experiments in Figure 4, which is an assumption that requires further testing). However, upon further inspection of the data, a comparison of Figures 1 and 4 shows that the hydrogen peroxide production rate in the experiment for the measurement of oxygen (see Figure 4) was more significantly reduced by the platinum electrode than it was in the experiments for the measurement of hydrogen (see Figure 1) (this can also be seen by comparing rows 2 and 4 of Table 1). The cause of this difference is not known: the only experimental condition that is different between the two experiments is the different purge gas that is used. Another effect of the purge gas was observed: over the course of a 1-h experiment, with helium purge gas, no significant change in either solution conductivity or pH was observed; however, when nitrogen was used as the purge gas, a decrease in pH (from ∼5.3 to 4.5) and an increase in solution conductivity (from ∼150 µS/cm to ∼162 µS/cm) were measured. It is not clear whether the changes in pH and conductivity and the change in hydrogen peroxide production are related, although since the basic problem is that there are more missing O atoms with nitrogen as the purge gas, the idea of the formation of small amounts of nitrogen oxides is inescapable (however, this was not measured). The numbers for the final hydrogen peroxide production rate for the oxygen measurement experiments are shown in the fourth row of Table 1. Whether the change in hydrogen production caused by the platinum electrode was the same in this case cannot be said, because a simultaneous measurement of the hydrogen and oxygen content was not made. In summary, there remain unanswered questions regarding the basic problem of the disparity between the effect of platinum on hydrogen on one hand, and hydrogen peroxide and oxygen collectively on the other hand, and that, to answer this question, simultaneous measurements of small concentrations of hydrogen
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Figure 5. Scanning electron microscopy (SEM) images of the two types of electrodes both before (top) and after (bottom) their use as high-voltage electrodes. The morphology of the “before” images is caused by the wire cutter, whereas the morphology of the “after” images is caused by the discharge.
and oxygen are needed. Potential techniques for this are being considered presently. However, the basic result, that the use of platinum as the high-voltage electrode significantly reduces the net production rates of all three species, can be stated categorically. 3.3. Location of the Heterogeneous Reaction. Because the only experimental difference between the two cases is the use of a different material as a high-voltage electrode, the effect must then be caused by a surface reaction, as discussed previously (excluding the remote possibility of an effect of dissolved metal ionssalthough nickel is known to undergo Fenton-type reactions, the nickel-chromium electrode is, here, the one that is apparently inert, compared to platinumsthe conductivity change discussed previously is probably not a result of dissolved metal ions, because it was apparently caused solely by the use of nitrogen as a carrier gas, suggesting the possibility that nitrogen oxides could be the cause of the change in pH and solution conductivity). It has been shown elsewhere that an electrical discharge in water from a metallic needle electrode causes a sputtering of metal particles into the water.12,13 Both types of needle electrodes in the present work were found to lose ∼2 mg of mass after 1 h of pulsed discharge in water. Therefore, there are two possible locations where the surface reactions may be occurring: the surface of the platinum needle itself or the surfaces of the solid platinum particles emitted from the electrode during the experiment. Figure 5 shows the effect of the pulsed discharge on the morphology of the needle electrodes. An attempt was made to retrieve and image a platinum particle that had been emitted from the electrode, and although an apparent group of particles was found with diameters in the range of ∼1-4.5 µm, further positive identification of these particles was not done. Back-of-the-envelope calculations show that the total surface area of 2 mg of particles of this size range would be on the order of a few square centimeters (assuming only 4.5-µm diameter particles gives ∼10 particles emitted per pulse, with a final total particle surface area of ∼1.4 cm2, and assuming only 1-µm-diameter particles gives ∼830 particles emitted per pulse, with a final total particle surface area of ∼5.6 cm2), compared with a surface area of only ∼8 mm2 for the needle electrode itself (1/32-in. diameter extended ∼3 mm into the solution). Thus, a simple comparison
of surface area suggests that the reactions likely occur on suspended particles and not the surface of the discharge needle itself. If platinum particles, which are continuously emitted into solution from the discharge, are responsible for the observed hydrogen peroxide and hydrogen data, it might be expected that the hydrogen and hydrogen peroxide concentrations and their corresponding net rates of production should continue to drop as the amount of platinum particles increases in solution. However, the net production rates of hydrogen and hydrogen peroxide are observed to not continuously drop toward zero, but to stabilize, as shown in Figure 3. A possible explanation for this fact is as follows. When the discharge in water is running, small bubbles can be visually observed to emanate from the discharge, swirl around through the solution (the reactor was stirred by a magnetic stirring bar) and eventually make their way to the gas/liquid interface and exit the solution, where the gas that they contained is then carried away to a gas sampling bottle by the purge gas (the purge gas bubbles were deliberately made not to pass close to the high-voltage needle, so as not to interfere with the discharge itself). These bubbles presumably contain hydrogen and oxygen in a ratio of ∼4:1, along with some water vapor (they are not steam bubbles, because any steam should condense upon moving away from the discharge). Because the sputtered particles are presumably present only in the solution, the small bubbles are essentially a vehicle for the gas to escape the reactor without coming into contact with the suspended particles. Therefore, the net production rate of hydrogen observed toward the end of an experiment, as seen in Figure 4, could simply be the amount of hydrogen that is able to escape the water in bubble form, and the degraded hydrogen, therefore must have dissolved from the bubbles into the solution and ended up reacting on the surface of the suspended particles. If this explanation is correct, then the stabilization of the net production rate of hydrogen peroxide may be explained also, because the measured hydrogen would not have been available to react with the remaining hydrogen peroxide. Another explanation might be that, because oxygen is also continuously introduced into the solution by the electrical discharge, it could be possible that the activity of the platinum particles for hydrogen and hydrogen peroxide degradation decreases, because of oxidation of the platinum by the molecular oxygen that is introduced into solution. At long times, there would then be a balance in the rates of formation of molecular species by the electrical discharge and the rates of degradation by surface reactions. Further work is necessary to analyze the particles emitted by the discharge and their chemical activity. While deactivation of the particles by oxygen is plausible, it should be noted that the platinum particles continuously emitted from the electrode are presumably active. It may also be noted that the consideration of irreversible adsorption of oxygen onto particles emitted by the discharge may also be applied to the simple case of the nickel-chromium electrode; in this case, there is “missing” oxygen as can be seen by observing Table 1 and Figure 4. In Table 1, the amount of measured oxygen is slightly less than what would be expected from the stoichiometry of reaction 1 (implying missing oxygen), and in Figure 4, the oxygen concentration takes a longer time to approach a stable value than does the hydrogen for the nickel-chromium electrode in Figure 1 (also implying missing oxygen). This “missing” oxygen could then have been taken up to form surface groups on the newly formed nickel-chromium particles.
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4. Conclusions Platinum reduces the production rates of hydrogen, hydrogen peroxide, and oxygen by electrical discharge in water. A set of reactions have been proposed to explain this result. An investigation was made into whether the responsible reactions occur on the surface of the discharge needle itself, or on platinum particles that are known to be emitted from the electrode. The findings suggest that these reactions occur on particles that are sputtered from the discharge electrode. Acknowledgment We would like to acknowledge the Florida Department of Environmental Protection for financial support. Thanks also to the FAMU-FSU College of Engineering, Department of Chemical Engineering, for support and to Wright Finney, Mayank Sahni, Petr Lukesˇ, and Selma Me{edovic´ for their discussion and assistance in the laboratory. We would also like to thank Kim Riddle (Biological Science Imaging Resource Group, Florida State University) for taking the electron micrographs. Literature Cited (1) Kirkpatrick, M. J.; Locke, B. R. Hydrogen, Oxygen, and Hydrogen Peroxide Formation in Aqueous Phase Pulsed Corona Electrical Discharge. Ind. Eng. Chem. Res. 2005, 44, 4243. (2) Lukes, P. Water Treatment by Pulsed Streamer Corona Discharge, Ph.D. Dissertation, Institute of Plasma Physics AS CR, Prague, Czech Republic, 2001. (ISBN: 80-902724-6-0.) (3) Joshi, A. A.; Locke, B. R.; Arce, P.; Finney, W. C. Formation of Hydroxyl Radicals, Hydrogen Peroxide and Aqueous Electrons by Pulsed Streamer Corona Discharge in Aqueous Solution. J. Hazard. Mater. 1995, 41, 3.
(4) Sato, M.; Ohgiyama, T.; Clements, S. Formation of Chemical Species and Their Effects on Microorganisms Using a Pulsed High-Voltage Discharge in Water. IEEE Trans. Ind. Appl. 1996, 32, 106. (5) Sun, B.; Sato, M.; Clements, J. S. Optical Study of Active Species Produced by a Pulsed Streamer Corona Discharge in Water. J. Electrost. 1997, 39, 189. (6) Sˇunka, P.; Babic´ky, V.; C ˇ lupek, M.; Lukesˇ, P.; Sˇimek, M.; Schmidt, J.; Cernak, M. Generation of Chemically Active Species by Electrical Discharges in Water. Plasma Sources Sci. Technol. 1999, 8, 258. (7) Sˇ unka, P. Pulse Electrical Discharges in Water and Their Applications. Phys. Plasmas 2001, 8, 2587. (8) Grymonpre, D. R.; Finney, W. C.; Clark, R. J.; Locke, B. R. Suspended Activated Carbon Particles and Ozone Formation in Aqueous Phase Pulsed Corona Discharge Reactors. Ind. Eng. Chem. Res. 2003, 42, 5117. (9) Sahni, M.; Finney, W. C.; Locke, B. R. Degradation of aqueous phase polychlorinated biphenyls (PCB) using pulsed corona discharges. In Proceedings of the Fourth International Symposium on Non-Thermal Plasma Technology for Pollution Control and Sustainable Energy DeVelopment, Panama City, FL, May 10-14, 2004. (10) Eisenberg, G. M. Colorimetric Determination of Hydrogen Peroxide. Ind. Eng. Chem. Res. 1943, 15, 327. (11) Huckaba, C. E.; Keyes, F. G. The accuracy of estimation of hydrogen peroxide by potassium permanganate titration. J. Am. Chem. Soc. 1948, 70, 1640. (12) C ˇ lupek, M.; Lukesˇ, P.; Babicky, V.; Sˇ unka, P.; Stefecka, M.; Skalny, J. D. Erosion of iron electrodes in pulsed corona discharge in water. In Book of Contributed Papers Part 2, 11th Symposium on Elementary Processes and Chemical Reactions in Low-Temperature Plasma, 1998. (13) Blokhin, V. I.; Vysikailo, F. I.; Dmitriev, K. I.; Efremov, N. M. Systems with Different Electrode Materials for Treatment of Water by a Pulsed Electric Discharge. High Temp. 1999, 37 (6), 963.
ReceiVed for reView October 14, 2005 ReVised manuscript receiVed January 23, 2006 Accepted January 31, 2006 IE0511480
