A Polarity Switching Technique for the Efficient Production of Sodium

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A Polarity Switching Technique for the Efficient Production of Sodium Hypochlorite from Aqueous Sodium Chloride Using Platinum Electrodes Rajkumar S. Patil, Vinay A. Juvekar,* and Vijay M. Naik Department of Chemical Engineering, Indian Institute of Technology Bombay, Powai, Mumbai-400 076, India ABSTRACT: A novel, periodic polarity-switching technique is developed for the electrolytic production of sodium hypochlorite from aqueous sodium chloride. The technique employs a flow-through cell having two identical platinum electrodes. Polarities of these electrodes are periodically swapped. The main purpose of this step is to prevent the growth of platinum oxide film and thereby retain high initial activity of the electrodes throughout the operation. The performance of this mode of operation is evaluated using the cumulative productivity and current efficiency of the cell as indicators. Typical ranges of parameters used are as follows: concentration of Cl− ions, 0.1−0.5 M; switching frequency, 0.01−1 Hz; and current density, 100−700 A m−2. The productivity of the switching mode is 5−6 times higher and its current efficiency is 15%−20% higher than that of the DC mode. The switching frequency has a strong influence on both the cumulative production rate and the average current efficiency and an optimal switching frequency exists under the given operating conditions. At high frequencies, the major side reaction is oxidation of the adsorbed hydrogen. As the concentration of NaCl increases, the optimal switching frequency decreases, and both the cumulative productivity and the average current efficiency increase. At the optimal frequency, the current efficiency of the switching mode varies from 60% to 95% as the concentration of NaCl is varied from 0.1 M to 0.5 M. The higher rate of circulation of the electrolyte improves the cell performance. The main cause for the loss in the current efficiency at high conversions of Cl− ions is the reduction of the hypochlorite ions at the cathode, which occurs under diffusion limiting conditions.

1. INTRODUCTION Sodium hypochlorite is a very strong oxidizing agent and has numerous applications, including treatment of drinking water,1 swimming pool water,2 and industrial wastewater; washing of textiles3/stained or soiled fabrics; sterilization of food processing equipment, surgical instruments and packaging materials; etc. Sodium hypochlorite, used in these applications, can be either manufactured/prepared in bulk or generated in situ at the point of use. In situ/onsite production of sodium hypochlorite can be achieved either by controlled reaction between already-prepared aqueous sodium hydroxide solution and chlorine or by electrolysis of sodium chloride solution. Of these methods, in situ electrolytic production has several advantages for the applications that require small quantities of sodium hypochlorite. These advantages include better safety due to elimination of the need to store and handle corrosive chemicals such as chlorine and sodium hypochlorite; superior dosage control; operational convenience; ready availability of the precursor (NaCl), etc. Generally, dimensionally stable anodes (DSAs) are employed for large-scale electrolytic production of sodium hypochlorite, because of their higher activity toward the oxidation of chloride (Cl−) ion. These anodes are prepared by applying a micrometer-thick coating of iridium oxide, ruthenium oxide, or mixed oxides on a titanium substrate.4 Platinum is an alternative electrode material. It is also known that freshly activated platinum anode shows much higher current efficiency than DSA.5,6 However, platinum electrode loses its initial activity/current efficiency quickly and, hence, it is less current efficient, compared to DSA, in longer run. The challenge lies in preserving and harnessing the better initial activity and current © XXXX American Chemical Society

efficiency of the platinum electrode over long operating periods of the electrochemical cell. The reason for the lower activity of platinum for the oxidation of NaCl is that, at the high anodic potentials, which are needed for generation of chlorine, platinum undergoes passivation due to formation of the oxide film. It has been observed that the passivation occurs in two stages.7,8 The first is a rapid-passivation stage, which is followed by a slowpassivation stage. During the first stage of passivation, the surface sites are progressively oxidized to platinum oxide. Since these oxide sites have much lower activity, compared to the metal sites, the activity of the electrode decreases as more and more metal sites are oxidized. At the later stages of operation, the rate of passivation slows, because of continuous renewal of the surface sites, as a result of the transport of the surface oxygen into the metal bulk by the place-exchange mechanism.9,10 After a sufficiently long time, the anode attains an almostconstant, albeit lower, value of the activity. The exchange current density on this passive electrode is substantially lower than that exhibited by a dimensionally stable electrode. On the other hand, activity of a freshly cathodically activated platinum electrode far exceeds that of DSA, as reported by Kraft et al.5 It is hypothesized that the following technique would improve the performance of an electrolytic cell producing NaOCl. The cell uses two platinum electrodes of identical size: one of them is the cathode and other is the anode. The polarity Received: August 1, 2014 Revised: November 21, 2014 Accepted: November 22, 2014

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dx.doi.org/10.1021/ie503084m | Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Industrial & Engineering Chemistry Research

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of the cell is switched periodically. During the first half of the cycle, the cathode not only performs the auxiliary reaction (reduction of water), but also undergoes self-activation, because of reduction of the oxide sites to metal. During the second half of the cycle, the cathode switches the polarity and becomes the anode. Since, this anode is now in the active state, it exhibits a high rate of oxidation of chloride ion. By optimizing the switching frequency, it would be possible to achieve cell performance, which is better than the cell that uses DSAs. Polarity switching in conjunction with platinum electrodes is also expected to be advantageous in the following respect. During electrolysis, hard ions from the feed solution deposit on the cathode (where pH is high), thereby drastically reducing the performance of the cell. It is known that intermittent reversal of the polarity of the electrodes helps in preventing this deposition.11,12 The reason is, when the cathode becomes the anode, the local pH at its surface reduces (due to generation of H+ ions). Most of the deposit dissolves at low pH and is flushed out of the cell. In the cells using DSA, intermittent switching of polarity results in corrosive depletion of the active coating material on the electrodes and necessitates its frequent replacement. The platinum is more resistant to corrosion, even under periodic switching;13 hence, it is expected to last much longer. The present work is aimed at studying the operating characteristics of a flow-through electrolytic cell having identical platinum electrodes in which the electrode polarity is periodically switched. Specifically, the productivity and current efficiency of the cell are studied as functions of the switching frequency. As a “control experiment”, the cell is also operated in DC mode, and the decay in its performance with time is monitored. The two modes of operation are compared to evaluate the advantage of the operation with periodic switching of polarity. The layout of the paper is as follows. We first list the possible reactions, which occur in the electrolytic cell and discuss their importance. We then describe the experimental setup and details of the electrolytic cell used in this study. This is followed by discussion on the effect of cell performance on various parameters such as concentration of sodium chloride, flow rate of the electrolyte through the cell, and polarity switching period. Next, we discuss the cathodic reduction of hypochlorite ions, which is an important side reaction, and we also quantify its effect on the current efficiency. Finally, we have compared the performance of the platinum electrode (operated under switching mode) with that of the DSA electrode.

2H 2O ⇔ O2 + 4H+ + 4e−

E 0 = 1.359 V

12ClO− + 6H 2O ⇔ 4ClO−3 + 12H+ + 8Cl− + 3O2 + 12e−

HOCl ⇔ ClO− + H+

(3)

pK a = 7.5

(5)

Pt + H 2O → Pt−OH + H+ + e−

(6a)

Pt−OH → PtO + H+ + e−

(6b)

At potentials above 1.2 V, platinum monoxide oxidizes further to dioxide: PtO + H 2O → PtO2 + 2H+ + e−

(7)

In a cell which is subjected to periodic switching, a significant anodic reaction is the oxidation of the adsorbed hydrogen. H 2(ads) → 2H+ + 2e−

E0 = 0 V

(8)

Hydrogen adsorption occurs on the electrode during the cathodic phase of the switching cycle. The major cathodic reaction is the reduction of water to hydrogen gas: 2H 2O + 2e− ⇔ H 2 ↑ + 2OH−

E 0 = −0.783 V

(9)

This reaction acts as the auxiliary reaction and provides the necessary current to sustain the anodic reactions. An undesirable cathodic reaction is the reduction of hypochlorite ions to chloride ions.

The chlorine generated in the reaction reacts with water to yield hypochlorite ions: (2)

E 0 = 0.46 V

This reaction has a much lower standard electrode potential, compared to that of chloride oxidation (i.e., reaction 1). However, its overpotential requirement appears to be high, since it is known that, the industrial chlorate cell, which works at NaOCl concentrations of 40−70 mM and uses anodes coated with noble metal (e.g., platinum, ruthenium, iridium), exhibits typical current efficiencies of 93%−95%. Moreover, the main side reaction governing the loss in current efficiency is the thermal decomposition of the hypochlorite ion to the Cl− ion and not its anodic oxidation.17 Therefore, for the present analysis, we have ignored this reaction. At the anodic potentials, platinum undergoes passivation, because of the formation of an oxide film. The oxide film can be both formed and reduced electrochemically.18−20 Breiter21 showed that most of the passivation occurs at very small oxygen coverage of the platinum surface. Gillman16 proposed the following mechanism for passivation, which postulates the formation of different oxides of platinum. At potentials of 0.2 mM). Hence, the range of measurement was limited to NaOCl concentrations of >0.2 mM. Electro-oxidation of Cl− ions is accompanied by many

4. RESULTS AND DISCUSSION 4.1. Cyclic Voltammogram of NaCl Solution at the Platinum Rotating Disk Electrode. Figure 3 shows cyclic voltammograms for three different concentrations of NaCl in the potential range of −2.5 V to +2.5 V (with reference to the Table 1. System and Material Data Used in These Studies: (a) Details of the Parallel Plate Electrolytic Cell, (b) Details of the Rotating Disk Electrode Cell, and (c) Properties of NaCl Solution at 298 K (a) Details of the Parallel Plate Electrolytic Cell parameter 1 2 3 4 5

value/comment

100 mm2 surface area of each electrode, Ae gap between electrodes, d 2 mm width of flow channel of cell, W 10 mm length of the channel, L 50 mm cell operating temperature 298 K (b) Details of the Rotating Disk Electrode Cell parameter

1 2 3 4 5 6 7

value/comment

type of cell single compartment diameter of rotating disk electrode 5 mm range of disk speed 100−10000 rpm reference electrode saturated calomel counter electrode Pt plate (10 mm × 10 mm) volume of the cell 200 mL cell operating temperature 298 K (c) Properties of NaCl solution at 298 K value/comment

1 2 3

parameter

0.1 M

viscosity of NaCl solutiona density of NaCl solutionb conductivity of NaCl solutionc

1.0 mPa s 1002 kg m−3 1.07 S m−1

0.25 M

0.5 M

1.0 mPa s 1.0 mPa s 1008 kg m−3 2.35 S m−1

1018 kg m−3 4.38 S m−1

1.0 M 1.0 mPa s 1040 kg m−3 8.44 S m−1

a

Viscosity of the solution is assumed to be the same as that of water. Estimated from density data from Perry’s Chemical Engineers’ Handbook.23 cMeasured by us.

b

D



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We also see that the hysteresis in the reverse cathodic sweep is very small and is likely to be caused by the double-layer capacitance. 4.2. Current Decay Characteristics of the Cell. Figure 4 presents a current density−time plot of the cyclic mode of

Figure 3. Cyclic voltammograms in different concentrations of sodium chloride (NaCl) at the platinum disk electrode. Scan rate, 0.3 V s−1; speed of disk, 3000 rpm; temperature, 298 K; volume of the solution, 150 mL; and pH, 6.7. The reference electrode is a saturated calomel electrode (SCE), the working electrode is a platinum rotating disk, and the counter electrode is a platinum wire mesh.

Figure 4. Cyclic mode of operation of the electrolytic flow cell. Cell voltage was switched between +2.81 V and −2.81 V every 5 s, and the current density across the electrodes was measured. The cell voltage is shown by the solid blue line and the current density is represented by the red line. NaCl solution (0.5 M) was circulated from a reservoir through the cell at a flow rate of 100 mL min−1; the volume of the solution in the reservoir was 50 mL, and the solution temperature was 298 K.

saturated calomel electrode, SCE). These voltammograms will be helpful in understanding the flow-cell performance. First, we find that, for all three concentrations of NaCl, the foot of the anodic wave lies at the same potential (1.13 V). This corresponds to the equilibrium potential for chloride oxidation reaction (reaction 1). We also observe that the current density varies linearly with electrode potential up to a certain potential. In the linear region, the platinum is very active and the current is solely controlled by the ohmic resistance of the solution surrounding the disk electrode.8 The inverse of the slope of the line represents the ohmic resistance. As the concentration of sodium chloride decreases, ohmic resistance increases and, hence, the slope of the line decreases, as seen from the figure. The electrode undergoes passivation at anodic potentials and when its surface activity falls to a sufficiently low value, the polarization curve exhibits a pseudo-plateau region, which is indicative of electrode passivation. As the concentration of NaCl decreases, the pseudo-plateau region occurs at lower potentials, indicating a greater rate of electrode passivation at lower concentrations of NaCl. When the electrode potential is increased well beyond the pseudo-plateau region, the current again begins to rise. In this region, the electrode surface is not the original platinum, but rather the platinum in its oxidized state. As a result of electrode passivation, marked hysteresis is observed during the reverse sweep. The anodic current is substantially lower than that during the forward sweep. Also, the foot of the reverse wave lies much to the right of the forward wave. A part of the cathodic region is enlarged in the inset of Figure 3. In the potential range between 0 and −0.6 V, there is a surface wave corresponding to hydrogen adsorption. Water reduction (reaction 9) begins at a potential of −0.85 V. Polarization curves in this region are the result of both ohmic resistance of the solution and the surface overpotential. We have not segregated the two effects. However, it is easy to observe that, for a given electrode potential, the current density increases as the concentration of Cl− ions increases.

operation of the flow cell. The plot was obtained after omitting the first few cycles, so that a cyclical steady state was achieved. The direction of the current undergoes reversal when the polarity is switched. Both the current and the potential in the reverse cycle were registered by the oscilloscope as negative quantities. It is also seen that the current density profiles of forward and reverse cycles have reflection symmetry, which indicates that both electrodes are practically identical in their behavior. At the beginning of a cycle, the current density is very high, but it undergoes a rapid initial decay followed by a slow decay later. The current efficiency during the initial decay phase was found to be low, indicating that one or more side reactions mainly contributed to the initial high current density. It is reasonable to assume that the major contribution is from the anodic oxidation of the adsorbed hydrogen gas (see reaction 8). Since the electrode undergoes a cyclical change in its polarity, it acts as a cathode during half of the cycle, during which, it adsorbs hydrogen generated by reduction of water. The electrode carries this hydrogen to the anodic phase of the cycle. The adsorbed hydrogen reduces the activity of the metal sites and must be oxidized to water in order to reactivate the sites. It is important to note that, in the cyclic mode, after the first transience, almost identical I−t curve was exhibited, cycle after cycle, until the end of the experiment. This indicates that the electrode is reactivated to the same state during every cycle and, hence, its activity does not decay with time. Figure 5 shows the decay of the current density during the DC mode. The anode was activated prior to the commencement of the run. This figure also shows rapid and slow currentdecay regimes. The initial region of the plot is enlarged in the inset of the figure. E



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mode, 20 cycles at equally spaced time points covering the entire time span of the experiment were selected and the area under the I−t plot for each of the selected cycle was computed using the trapezoidal rule. The arithmetic average of these 20 cycles was then used to compute the total charge by multiplying the average value of cycle charge by the number of cycles elapsed during the experiment. Figure 6 compares the productivity, P(t), of DC and cyclic modes at different values of the operating time (t).

Figure 5. Decay of current density with time during the DC mode of operation of the cell. [Parameters: NaCl concentration, 0.5 M; electrolyte flow rate, 100 mL min−1 (Re = 278); cell voltage, 2.81 V; volume of solution in the reservoir, 50 mL; and solution temperature, 298 K.]

The inset in Figure 5 shows that the rapid decay is not as sharp as that observed in the cyclic mode, possibly indicating that a greater amount of hydrogen is absorbed during the preactivation stage (probably due to the longer activation time of 300 s, which was used in this case). However, the current density at the end of the rapid decay regime is the same as that observed in the cyclic mode, indicating that the electrode attains the same state of activation after the hydrogen is completely oxidized. Comparison of Figures 4 and 5 shows that, in the cyclic mode, the cell exhibits much higher current density than the time-averaged current density exhibited by the batch mode. The reason is, although, in both modes, the initial activity of the anode is almost the same, cyclic mode does not allow the electrode activity to decay with time. 4.3. Cell Performance. Two relative indicators are used to evaluate the performance of the cell. The first is the cumulative productivity of the cell up to time t, per unit area of the electrode. It is denoted by P(t) (expressed in units of mol m−2) and computed using the following equation: P(t ) =

C NaOCl(t )νR Ae

Figure 6. Comparison of cumulative productivity of switching mode with DC mode. Both modes are operated with 0.5 M NaCl at a cell voltage of 2.81 V and an electrolyte flow rate of 100 mL min−1; the volume of solution in the reservoir is 100 mL in DC mode and 150 mL in switching mode, and the solution temperature was 298 K. A switching time of 10 s is used in cyclic mode, which is represented by the solid blue line; that in DC mode is shown by the dotted red line.

The cumulative productivity of the cyclic mode is seen to be substantially higher than that observed in the DC mode. The reason for this difference lies in periodic rejuvenation of the activity of the electrode in the cyclic mode, as explained in section 4.2. Figure 7 compares the current efficiencies of two modes. This figure shows that, up to a time period of ∼3000 s, the initial current efficiency of the cyclic mode is >95%, whereas that of the DC mode is only 73%. In both modes of operation,

(11)

where CNaOCl(t) represents the concentration of NaOCl in the reservoir at time t, νR is the volume of the electrolyte in the reservoir, and Ae is the area of the electrode. The second indicator is the cumulative percent current efficiency of the cell, ηC(t). It is the percentage of the consumed electric charge that is utilized in the production of NaOCl. It is obtained using the following equation: ηC (t ) =

2FC NaOCl(t )vR × 100 Q (t )

(12) Figure 7. Comparison between the cumulative current efficiency of cyclic and DC modes. Both modes are operated with 0.5 M NaCl at a cell voltage of 2.81 V, with an electrolyte flow rate of 100 mL min−1; the volume of solution in the reservoir was 50 mL, and the solution temperature was 298 K. A switching time of 10 s is used in cyclic mode, which is represented by the solid blue line; that used in DC mode is shown by a dotted red line.

where F is the Faraday’s constant and the factor of 2 takes into account the fact that two electrons must be removed at the anode in order to produce one molecule of NaOCl. Q(t) is the quantity of electric charge consumed by the cell up to time t. For the batch mode, Q(t) was estimated by computing the area under the I−t plot, using the trapezoidal rule. For the cyclic F



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the cumulative current efficiency decreases with time. It will be shown in section 4.7 that the reduction in the current efficiency in the cyclic mode is mainly caused by the cathodic reduction of hypochlorite ions (recall reaction 10). 4.4. Effect of Switching Period. Figure 8 presents the effect of the switching period on cumulative productivity and current efficiency of the cyclic mode for 900 s of operation of the cell.

Figure 9. Effect of Reynolds number (Re) on the cumulative productivity of the cell. The cell is operated for 900 s with 0.5 M NaCl, at a cell voltage of 2.81 V. The temperature of the electrolyte was 298 K, the volume of solution in the reservoir was 50 mL, and the polarity switching time was 10 s.

Figure 8. Effect of polarity switching period on cumulative productivity and current efficiency. The cell is operated for 900 s with 0.5 M NaCl, at a cell voltage of 2.81 V and an electrolyte flow rate of 100 mL min−1. The volume of solution in the reservoir was 50 mL, and the solution temperature was 298 K. The solid green line represents the productivity, and the dotted purple line represents the current efficiency.

Both P(900) and ηC(900) peak at almost the same switching period, indicating the existence of an optimal switching frequency. At very short switching periods, oxidation of the adsorbed hydrogen is the dominant side reaction, and, hence, the performance is poor. The performance picks up as hydrogen is consumed and active electrode surface is made available. However, passivation of the anode becomes significant as the switching period increases. Passivation not only causes a reduction in the rate of chloride oxidation, but also promotes the oxidation of water. Current efficiency (ηC(t)) exhibits a very gradual decline beyond the peak, compared toP(t). 4.5. Effect of Rate of Circulation of Electrolyte. Figure 9 shows the cumulative productivity of the cell against the Reynolds number (Re) for the flow of electrolyte through the cell. There is a pronounced effect of Re on the performance of the cell. The productivity increases sharply with Re in the lower range of Re values, but attains a plateau at high Re. This result is surprising, since the rate of oxidation of NaCl is not controlled by mass transfer. To probe into this phenomenon further, in Figure 10, we have plotted the current density in the cell as a function of time in a cycle, using the Reynolds number (Re) as the parameter. We see that the current density increases as the circulation velocity increases. Since the productivity is directly proportional to current density, the trend in Figure 10 is consistent with that in Figure 9. We wish to point out that, under periodic switching conditions, the anode remains very active and offers negligible overpotential for charge transfer. Hence, the current density is governed by the rate of the cathodic reaction and the ohmic resistance of the cell. At high current densities prevailing in the cell, hydrogen can nucleate in the form of gas bubbles (swarms

Figure 10. Effect of electrolyte flow rate on current decay characteristics with time. The cell is operated with 0.5 M NaCl, at a cell voltage of 2.81 V; the volume of solution in the reservoir was 50 mL, the solution temperature was 298 K, and the switching time was 10 s. The trace of current versus time is obtained in an arbitrarily chosen cycle after cyclical steady state is attained; t = 0 corresponds to the time at which the polarity is switched. Different lines correspond to different Re values, as listed here: solid black line, Re = 27.8; red line with triangles, Re = 83.4; blue line with squares, Re = 111; green line with circles, Re = 167; and violet line with stars, Re = 278. The inset shows a I−t plot where the shaded area represents the contribution by the charge needed for hydrogen oxidation, and the remainder of the area under the curve represents the contribution by oxidation of the Cl− ion.

of gas bubbles were visually observed to be issued from the gap between the electrodes). These bubbles had a tendency to stick to the metal surface, as was confirmed visually. They can cover part of the surface area of the cathode and thereby hinder the charge transfer reaction. The circulating liquid disengages hydrogen gas bubbles and thereby increases the available cathodic area for charge transfer. The higher the rate of circulation, the more efficient the bubble disengagement, which allows more cathodic reaction to occur. At a sufficiently high circulation rate, the drag is high enough to completely prevent sticking of bubbles. This limit is seen to be attained at Re ≈ 200 G



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in Figure 9. The current density increases slowly beyond this velocity. 4.6. Effect of NaCl Concentration. Figure 11 shows the effect of NaCl concentration on the cumulative productivity at

Figure 12. Effect of NaCl concentration on the cumulative current efficiency of the cell. The cell is operated for 900 s in switching mode at a cell voltage of 2.81 V; the temperature of the electrolyte is 298 K, the volume of solution in the reservoir was 50 mL, the solution temperature was 298 K, and the electrolyte flow rate was 100 mL min−1 (Re = 278). The solid blue line corresponds to 0.5 M NaCl, the dashed green line corresponds to 0.25 M NaCl, and the dotted red line corresponds to 0.1 M NaCl.

Figure 11. Effect of NaCl concentration on the cumulative productivity of the cell. The cell is operated for 900 s at a cell voltage of 2.81 V and an electrolyte flow rate of 100 mL min−1; the volume of solution in the reservoir was 50 mL, and the solution temperature was 298 K. The solid blue line corresponds to 0.5 M NaCl, the dashed green line corresponds to 0.25 M NaCl, and the dotted red line corresponds to 0.1 M NaCl.

a constant cell voltage at 2.81 V. This figure shows that the productivity of the cell increases with NaCl concentration. We have seen from Figure 3 that the values of equilibrium potential are 1.13 V for anodic reactions and −0.85 V for cathodic reactions. The total driving force for the current equals [2.81 V − (1.13 V − (−0.85 V))] = 0.83 V. This figure also shows that, on an active platinum electrode, the overpotential for the anodic reaction is negligibly small. Hence, the potential driving force of 0.83 V would be consumed partially by the ohmic drop and partially by the cathodic overpotential. From the cathodic polarization curves in Figure 3, we have seen that, for the same electrode potential, current density at the cathode increases as the NaCl concentration increases. On this basis, we expect the cell current to increase as the NaCl concentration increases for a fixed cell voltage. Consequently, greater productivity is expected at higher NaCl concentrations. In Figure 12, the effect of NaCl concentration on current efficiency at different switching periods is presented. It is wellknown that, with increasing NaCl concentration, the electrode is less susceptible to passivation. This is reflected in the effect of switching period on the cumulative percent current efficiency. In the case of 0.1 M NaCl, there is a substantial decrease in the current efficiency with increases in the switching period, whereas in 0.5 M NaCl, very little passivation is observed, even when the switching period is increased up to 50 s. 4.7. Cathodic Reduction of Hypochlorite. Sodium hypochlorite undergoes cathodic reduction according to reaction 10. This reaction adversely affects the cell performance. In order to determine the kinetics of this reaction, linear sweep voltammetry was conducted on a rotating disk electrode in 1 M NaCl solution containing different concentrations of NaOCl. The potential was swept from 0.9 V to −1.25 V at a scan rate of 0.5 V s−1. Figure 13 shows the voltammograms in the potential range of interest. Two plateaus are seen in Figure 13. The first lies in the potential range from 0 to −0.2 V, and the second appears

Figure 13. Linear sweep voltammograms of NaOCl solution by using RDE. [Parameters: Disk speed, 3000 rpm; scan rate, 0.5 V s−1; NaCl concentration, 1.0 M; electrolyte temperature, 298 K; electrolyte volume, 100 mL; working electrode, platinum disk; counter electrode, platinum foil; and reference electrode, SCE.] Line colors correspond to different concentrations of NaOCl prepared in 1.0 M NaCl: solid black line, 0 mM; dashed red line, 4.99 mM; dotted blue line, 9.52 mM; dash-dotted green line, 13.7 mM; dash-dot-dotted magenta line, 17.5 mM; and closely dotted brown line, 21.0 mM.

beyond −0.6 V. The first plateau is formed due to overlap of the rising reduction wave of hypochlorite ions and the decreasing part of the hydrogen adsorption wave (the surface wave for hydrogen adsorption lies between 0 and −0.6 V, as seen from Figure 3). The second plateau corresponds to the limiting current for hypochlorite reduction. The plateau continues until the water reduction reaction (reaction 9) becomes predominant at electrode potentials beyond approximately −1.0 V. For the purpose of estimating the limiting current density, iL, we have used the region enclosed between the two dotted lines shown in Figure 13. In Figure 14, we have plotted iL against the concentration of sodium hypochlorite in the solution. H



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In this equation, the left-hand side is the net rate of accumulation of sodium hypochlorite in the system. The term rc on the right side represents the rate of generation of hypochlorite per unit area of the anode, and the second term represents its consumption due to reduction at the cathode. The rate of generation (rc) is assumed to be independent of time under cyclic mode, since the electrode activity remains constant. Also, only a small fraction of the Cl− ions in the reservoir are consumed, so the NaCl concentration can also be considered to be constant. Equation 14 can be integrated under the initial condition c(0) = 0, to yield an expression for the concentration of NaOCl as a function of time: C NaOCl(t ) =

Figure 14. Limiting current density versus NaOCl concentration. The slope of the straight line is 20.82 A mol−1 m−1.

⎡ ⎛ t ⎞⎤ P(t ) = rcτd⎢1 − exp⎜ − ⎟⎥ ⎢⎣ τd ⎠⎥⎦ ⎝

(16)

where τd, which represents the time constant for the cathodic reduction of sodium hypochlorite, is defined as v τd = R kLAe (17)

where n is the number of electron transferred (here, n = 2); D is the diffusion coefficient (expressed in units of m2 s−1); ν is the kinematic viscosity of electrolyte solution (m2 s−1); F is the Faraday’s constant (C mol−1); ω is the angular velocity of the disk electrode (rad s−1); and CNaOCl is the bulk concentration of hypochlorite ion (mol m−3). The slope of the line is 20.8 A mol−1 m−1, from which we can obtain the diffusion coefficient of hypochlorite ion as D = 1.63 × 10−9 m2 s−1. We can justify this value by comparing it with diffusion coefficients at infinite dilution at 25 °C, of two analogous monovalent ions,24 i.e., the chlorate ion (D = 1.719 × 10−9 m2 s−1) and the perchlorate ion (D = 1.792 × 10−9 m2 s−1). After allowing for the fact that the diffusion coefficient in finite electrolyte concentration (1 M in this case) is lower than that at infinite dilution, we can justify the estimated value of hypochlorite ion. Since (i) the reduction of sodium hypochlorite yields sodium chloride as the only product and (ii) sodium chloride is already present in the system in large excess, the extent of hypochlorite reduction in the cell cannot be directly sensed. It can only be indirectly quantified through the decrease in the net rate of production of hypochlorite with time. In the cell, the reduction of hypochlorite ions is expected to occur under the limiting current conditions. The reason is that the major cathodic reaction is the reduction of water, which occurs at a potential of approximately −1.0 V (with respect to SCE); this potential is much greater than that at which the hypochlorite reduction reaction attains the limiting current condition, as seen from Figure 13. It is thus possible to estimate the rate of reduction of hypochlorite ion in the cell and use this information to quantify the effect of the reaction on the cell performance. Since the trend of productivity, shown in Figure 6, reflects this effect, we have used this figure for the analysis. Consider first the cell operation in the periodic mode, which corresponds to the upper curve in Figure 6. The mass balance of sodium hypochlorite in the system can be represented by the following equation: dC NaOCl(t ) = rcAe − kLAeC NaOCl(t ) dt

(15)

Using eq 11, we can convert eq 15, in terms of the cumulative production of sodium hypochlorite:

The plot is a straight line passing through the origin. Using the slope of the line, it is possible to calculate the diffusion coefficient of hypochlorite ion using the Levich equation:25 iL = 0.62nFD2/3(ν)−1/6 ω1/2 C NaOCl (13)

vR

⎛ k A t ⎞⎤ rc ⎡ ⎢1 − exp⎜ − L e ⎟⎥ kL ⎢⎣ vR ⎠⎥⎦ ⎝

We have fitted the data of the cyclic mode of Figure 6 to eq 16. In this case, the NaCl concentration is considered to be constant in this analysis (the conversion of NaCl during electrolysis is estimated to be only 3.1% over a total period of operation of 104 s). The rate rc is computed from the current density decay profile in the fresh NaCl, where the only reactions are NaOCl generation and oxidation of adsorbed hydrogen. The total charge consumed during the half cycle (Qc) is obtained by integration of the current density−time (I− t) plot and the charge corresponding to hydrogen oxidation (Qh) was subtracted from it. The charge Qh was estimated by measuring the shaded area below the curve, as shown in the inset in Figure 10. The shaded area is located above the tangent drawn to the curve at a point where the rapidly decaying current (caused by rapid oxidation of adsorbed hydrogen) meets the slowly decaying current (caused by electrode passivation). We assume that the unshaded area has no contribution from hydrogen oxidation. The following equation is used to estimate rc: rc =

Qc − Qh 2Fτ

(18)

where τ is the switching time. For the data in Figure 6 (upper curve), τ = 10 s, Qc = 0.429 C, and Qh = 0.0129 C. The value of rc is estimated to be 2.16 × 10−3 mol m−2 s−1. Since the contribution to the total charge due to hydrogen oxidation is ∼10% at the lowest circulation rate, the maximum error in the procedure outlined here is expected to be very small. From eq 16, we expect the plot of −ln(1 − P(t)/ri) versus time to be a straight line passing through the origin and having a slope of 1/τd. This plot, for the best fit value of τd = 1.26 × 104 s, is shown in Figure 15 (filled circles). The fit is excellent. The value of kL, as estimated from τd, using eq 17, is 1.19 × 10−4 m s−1. It is worthwhile to compare this value with that obtained using a Leveque-type

(14) I



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We have used eq 21 to estimate both rc and τd from the data on DC mode in Figure 6. The initial point is taken as the first point in the figure, for which t0 = 900 s and P(t0) = 0.7198 mol m−2. We expect the electrode to attain almost constant activity by this time, based on the trend observed in Figure 5. The bestfit values are rc = 7.24 × 10−4 mol m−2 s−1 and τd = 9.77 × 103 s. The fit of the model to the experimental data is shown in Figure 15. This fit is also very good. Comparing the values of rc in the two modes, we conclude that the rate of production of NaOCl in the cyclic mode is ∼3.5 times greater than that in the DC mode. This difference is caused both by lower activity and lower current efficiency of the DC mode. The mass-transfer coefficient in the DC mode is estimated from τd = 9.77 × 103 s as kL = 1.02 × 10−4 m s−1, which is lower than that observed in the cyclic mode (kL = 1.19 × 10−4 m s−1). This is understandable, since the rate of hydrogen evolution in the DC mode is lower than that in the cyclic mode, because of the lower current density in the former. 4.8. Comparison with DSA. It is important to compare the switching mode of operation using platinum electrodes and DC mode operation using DSA. Comparison of the current efficiency is made as shown in Figure 16. In this figure, the

Figure 15. Plot of −ln(1 − P(t)/ri) versus time t: solid red squares (■) denote DC mode data, and solid blue circles (●) represent switching mode data.

correlation,26 which is valid for mass transfer in a parallel plate cell under fully developed laminar flow: ⎛ ReScde ⎞1/3⎛ 2 ⎞ ⎟ ⎜ Sh = 1.467⎜ ⎟ ⎝ L ⎠ ⎝ γ + 1⎠

1/3

(19)

In this equation, Sh is the Sherwood number (Sh = kLde/D), Sc is the Schmidt number (Sc = ν/D), L is the length of the flow channel, and γ = d/W (where W is the width of the channel and d is its depth). The value of kL, as estimated using eq 19, is 1.30 × 10−5 m −1 s , which is about one-ninth of the observed value of 1.19 × 10−4 m s−1. We reconcile this difference using the fact that hydrogen evolution at the cathode enhances the mass-transfer coefficient. The effect of gas bubbles on enhancement of the mass-transfer coefficient has been reviewed by Wendt and Kreysa.27 They have provided the following correlation for approximate estimation of mass-transfer coefficients in gasevolving electrodes: Sh = 0.93(ReSc)0.5

(20)

This equation predicts a mass-transfer coefficient value of 1.87 × 10−4 m s−1, under the conditions of Figure 6, which is of the same order of magnitude of the observed mass-transfer coefficient. It would be interesting to extend the same analysis to the DC mode of operation. A problem in analyzing this mode is that the electrode activity continuously decreases with time. Hence, in this mode, rc cannot be considered to be constant, at least during the initial period. However, after a certain lapse of time, electrode activity decreases very slowly (as seen from Figure 5) and can be considered to be approximately constant. Equation 15 can be applied beyond this point in time. Let this time be denoted by t0 and the corresponding concentration of NaOCl be denoted by CNaOCl(t0). We solve eq 15 using the condition that c(t) = c(t0) at t = t0 and obtain the following equation, in terms of cumulative productivity: ⎡ ⎛ ⎛ t − t 0 ⎞⎤ P(t0) ⎞ P(t ) = rcτd⎢1 − ⎜1 − ⎟ exp⎜ − ⎟⎥ ⎢⎣ rcτd ⎠ τd ⎠⎥⎦ ⎝ ⎝

Figure 16. Comparison of the cumulative current efficiency of the IrO2 electrode in DC mode with the Pt electrode in switching mode. [Legend: solid line (), IrO2 coated on titanium expanded metal electrodes operated at a chloride concentration of 19 g L−1 at a current density of 15 mA cm−2 and a temperature of 23 °C; solid square (■), Pt electrode.]

solid line represents the current efficiency of the iridium oxide anode, as observed by Kraft et al.2 The data points represent our data on platinum, relative to the switching mode of operation. It is seen that the platinum electrode in the switching mode is better than the IrO 2 electrode when high concentrations of NaCl solution (>15 g dm−3) are used. Kraft et al.5 have also compared the active chlorine production rate on a freshly activated platinum electrode with that on an IrO2 electrode under identical operating conditions. The rate of production on the freshly activated platinum electrode is 37 mg A−1 h−1. On the IrO2 electrode, the rate of production at the steady state (where its activity is highest) is 22 mg A−1 h−1. Since the platinum electrode, operated in cyclic mode, should resemble the freshly activated electrode, in terms of its performance, we can conclude that it would be ∼50% more active than the IrO2 electrode.

(21)

where P(t0) = c(t0)νR/Ae. Hence, the plot of ln[1 − P(t)/ (rcτd)] versus time should be a straight having a slope of 1/τd and the intercept on the ordinate should be −ln[1 − P(t0)/ (rcτd)] − t0/τd. J



Article

5. CONCLUSIONS Through this study, we have demonstrated the advantage of using a periodic polarity-switching technique involving platinum electrodes for the electrosynthesis of NaOCl. Both the cumulative production rate and the current efficiency show significant increases over the DC mode of operation. The active platinum electrode is also better than DSA, with respect to both productivity and the current efficiency. There is an optimum switching frequency that balances the low current efficiency at short switching periods, which is due to the hydrogen oxidation reaction, and the lower electrode activity, which is due to passivation at longer switching periods. A high NaCl concentration is preferred, since, in this case, not only does the cell have a higher productivity, but electrodes also are less susceptible to passivation. To achieve high cell efficiency, a NaCl concentration of 0.5 M (or higher) is advisible. A high rate of circulation has a beneficial effect on cell performance. In a cell with active platinum electrodes and a high circulation rate of electrolytes, the only controlling resistance is the ohmic resistance of the cell. As the concentration of NaOCl in the system increases, the cathodic reduction of NaOCl becomes significant. The reduction of NaOCl occurs under limiting current conditions. Bubbles of hydrogen, which are generated at the cathode, enhance the rate of reduction of hypochlorite by enhancing the rate of mass transport.

■

rc = rate of generation of hypochlorite ion per unit area of anode, mol m−2 s−1 Sc = Schmidt number Sh = Sherwood number T = temperature, K t = time, s W = width of the channel, m Greek Symbols

■

ηC(t) = cumulative percent current efficiency λmax = wavelength of light at which maximum absorption is observed, nm ν = kinematic viscosity of the solution, m2 s−1 ω = angular velocity of the rotating disk electrode, rad s−1 τ = switching time, s τd = time constant for cathodic reduction of sodium hypochlorite (defined by eq 17), s γ = ratio of the depth to width of flow channel; γ = d/W

REFERENCES

(1) White, G. C. Handbook of Chlorination; Van Nostrand Reinhold Company: New York, 1986. (2) Kraft, A.; Stadelmann, M.; Blaschke, M.; Kreysig, D.; Sandt, B.; Schroder, F.; Rennau, J. Electrochemical water disinfection Part I: Hypochlorite production from very dilute chloride solutions. J. Appl. Electrochem. 1999, 29, 861−868. (3) Szpyrkowicz, L.; Juzzolino, C.; Kaul, S. N. A comparative study on oxidation of disperse dyes by electrochemical process, ozone, hypochlorite and fenton reagent. Water Res. 2001, 35, 2129−2136. (4) Beer, B. H.; Schiedam, N.; , Electrode and method of making same. U.S. Patent 3,234,110, 1966. (5) Kraft, A.; Blaschke, M.; Kreysig, D.; Sandt, B.; Schroder, F.; Rennau, J. Electrochemical water disinfection Part II: Hypochlorite production from potable water, chlorine consumption and the problem of calcareous deposits. J. Appl. Electrochem. 1999, 29, 895− 902. (6) Tilak, B. V. Kinetics of chlorine evolutionA comparative study. J. Electrochem. Soc. 1979, 126, 1343−1348. (7) Kuhn, A. T.; Wright, P. M. A study of the passivation of bright platinum electrodes during chlorine evolution from concentrated sodium chloride solutions. J. Electroanal. Chem. Interfacial Electrochem. 1972, 38, 291−311. (8) Patil, R. S.; Juvekar, V. A.; Naik, V. M. Oxidation of chloride ion on platinum electrode: dynamics of electrode passivation and its effect on oxidation kinetics. Ind. Eng. Chem. Res. 2011, 50, 12946−12959. (9) Sato, N.; Notoya, T. Measurement of the anodic oxide film growth on iron for hours. J. Electrochem. Soc. 1967, 114, 585−586. (10) Sato, N.; Cohen, M. The kinetics of anodic oxidation of iron in neutral solution. J. Electrochem. Soc. 1964, 111, 512−519. (11) James, A. T.; Brian, R. H.; Ajit, K. C.; Jeremy, J. V.; Dennis, E. B.; Myron, F. M.; Karl, W. M. Method and apparatus for the electrochemical treatment of liquids using frequent polarity reversal. U.S. Patent 2011/0079520A1, 2011. (12) Justin, A. A.; Kevin, S. A.; Rodney, E. H. Reverse polarity cleaning and electronic flow control systems for flow intervention electrolytic chemical generation. U.S. Patent 2009/0229992A1, 2009. (13) Kraft, A. Electrochemical water disinfection: A short review. Platinum Met. Rev. 2008, 52, 177−185. (14) Obrucheva, D. The platinum electrode X. A study of the adsorption of oxygen by smooth platinum by the electrochemical method. Zh. Fiz. Khim. 1952, 26, 1448−1457. (15) Breiter, M. W. Voltammetric study of halide ion adsorption on platinum in perchloric acid solutions. Electrochim. Acta 1963, 8, 925− 935. (16) Gilman, S. Electrochemical surface oxidation of platinum. Electrochim. Acta 1964, 9, 1025−1046.

AUTHOR INFORMATION

Corresponding Author

*Tel.: +91-22-25767236. Fax: +91-22-25766895. E-mail: vaj@ iitb.ac.in. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS The authors wish to acknowledge Unilever Industries Private Limited for funding this research.

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LIST OF NOMENCLATURE Ae = area of electrode, m2 CNaOCl = concentration of sodium hypochlorite, M CNaCl = concentration of sodium chloride, M D = diffusion coefficient, m2 s−1 d = gap between two electrode (depth of flow channel), m de = hydrodynamic equivalent diameter, m E = electrode potential, V Ecell = cell voltage, V E0 = standard electrode potential, V F = Faraday constant, C mol−1 i = current density, A m−2 iL = diffusion limiting current density, A m−2 k = electrical conductivity of NaCl solution, S m−1 kL = mass-transfer coefficient, m s−1 n = number of electron transferred P(t) = cumulative productivity, mol m−2 s−1 Q(t) = quantity of electric charge, C Qc = quantity of electric charge consumed during first half cycle, C Qh = quantity of electric charge due to hydrogen oxidation, C R = universal gas constant, J mol−1 K−1 Re = Reynolds number K



Article

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A Polarity Switching Technique for the Efficient Production of Sodium

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