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Electrodissolution of Aluminum Electrodes in Electrocoagulation Processes P. Can ˜ izares, M. Carmona, J. Lobato, F. Martı´nez, and M. A. Rodrigo* Department of Chemical Engineering, Facultad de Ciencias Quı´micas, Universidad de Castilla La Mancha, Campus Universitario s/n, 13071 Ciudad Real, Spain
The electrodissolution of aluminum electrodes in aqueous solutions containing sulfate or chloride ions is studied in this work. The results obtained are important in order to obtain a better understanding of the electrocoagulation process, as the electrodissolution of the anode surface is its first step. It has been determined that both chemical and electrochemical dissolution play an important role in the aluminum generation. The chemical dissolution of aluminum is strongly influenced by the pH. Alkaline pHs increase the dissolution rate by orders of magnitude. Within the experimental conditions used, the supporting media does not seem to influence greatly the chemical dissolution process. The electrochemical dissolution process depends mainly on the specific electrical charge passed. Salinity does not significantly affect the electrodissolution rate. Good fittings between experimental and modeled data are obtained by modeling the system with a simple model based on two assumptions: a highly segregated flow pattern and the calculation of aluminium species and pH from a pseudoequilibrium approach. Introduction During the last few decades, electrochemical wastewater treatment technologies have undergone rapid development. One of these technologies is the electrochemically assisted coagulation that can compete with the conventional chemical coagulation process in the treatment of wastes polluted with colloids or macromolecules or in the treatment of emulsions.1-8 Electrochemically assisted coagulation was seen as a promising technology at the turn of the nineteenth century, although it has received very little scientific attention. In the following decades, plants were commissioned in the United States to treat municipal wastewater. However, all such plants were abandoned due to apparent higher operating costs and some expectations of high initial capital costs as compared to the case of chemical dosing. Recently, some works focused on the study of this technology have reported promising results, and several firms have started to commercialize plants for treating actual wastes. Nevertheless, to make this technology competitive with the conventional chemical coagulation process, more effort has to be done in the next years, and a better understanding of the processes involved must be achieved. The main stages involved in the electrochemically assisted coagulation are shown schematically in Figure 1.9-10 An electrochemical cell is used to provide aluminum or iron to the waste. These coagulant reagents desestabilize the colloidal pollutants (or break the emulsion), similar to the case of a conventional coagulation process. The turbulence generated by the oxygen and the hydrogen evolution generates a soft mix that helps the desestabilized colloids to flocculate (to link together and to generate bigger particles). This process * To whom correspondence should be addressed. Tel.: +34 902 20 41 00. Fax: +34 926 29 53 18. E-mail:
[email protected].
Figure 1. Main stages involved in the electrochemically assisted coagulation.
is often called electroflocculation. Finally, the pollutants are removed by sedimentation, filtration, or flotation from the waste. The bubbles formed in the oxygen and mainly in the hydrogen evolution can help to increase the efficiency of the later process (electroflotation).11-16 Most works in the literature consider that the coagulant is obtained mainly by the electrodissolution of the anode surface, but recently, some authors17 have proposed that, to justify the electrocoagulation results, the chemical dissolution of the electrodes should also be considered in this representation of the electrocoagulation process. The aim of this work is to study the electrodissolution of the aluminum into wastes. The characterization of this first step is important in order to have a complete understanding of the electrochemical coagulation process. To do this, the chemical dissolution of the aluminum sheets and the electrochemical dissolution has been studied and a model that considers both chemical and electrochemical processes is proposed.
10.1021/ie048858a CCC: $30.25 © 2005 American Chemical Society Published on Web 05/13/2005
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Figure 2. Layout of the bench-scale plant, including detail of the electrochemical section.
Experimental Section Experimental Devices. The electrochemical dissolution experiments have been carried out in a continuous bench-scale plant, with a single-compartment electrochemical flow cell (Figure 2). Aluminum electrodes (HE 18) were used as the anode and cathode. Both electrodes were square in shape (100 mm sides) with a geometric area of 100 cm2 each and with an electrode gap of 9 mm. The electrical current was applied using a DC Power Supply, FA-376 Promax. The current flowing through the cell was measured with a Keithley 2000 Digital Multimeter. The electrolyte was stored in a 5000 mL glass tank stirred by an overhead stainless steel rod stirrer, Heidolph RZR 2041, and thermostatized by means of a water bath to maintain the temperature at the desired set point, and circulated through the electrolytic cell by a peristaltic pump. Chemical dissolution experiments were carried out in stirred batch reactors. To do this, a multistirrer device, Isco, was used. Experimental Procedure. Electrochemical dissolution experiments were carried out in the continuousoperation bench-scale plant, under galvanostatic conditions. The electrolyte was pumped into the cell and collected in a tank. Samples were taken at the outlet of the cell, and pH was measured using an inoLab WTW pH meter. To determine the total aluminum concentration, samples were diluted 50:50 v/v with 4 N HNO3, and after that, they were measured using a sequential inductively coupled plasma spectrometer, Varian Liberty, according to an standard method18 (plasma emission spectroscopy). Preparative experiments were carried out to determine the amount of passed solution required to reach the steady state. In every experiment, this final state was tested. To determine the influence of the passed electric current charge (I/Q), the flow rate (Q) was maintained constant at 19.8 dm3 h-1 and the employed current density was changed. The range of the passed electrical charge studied was 0.00070.0138 A h dm-3. To determine the influence of the pH, the feeding solution pH was changed in the range 1-13. To determine the influence of the current density, the current density and the flow rate were changed simultaneously in order to obtain the same amount of passed current charge. To test the influence of the supporting
Figure 3. Aluminum dissolved and pH variation during batch chemical dissolution experiments carried out under nonalkaline conditions: [ 2450 mg dm-3 NaCl, pH 3; 0 3000 mg dm-3 Na2SO4, pH 3; 2 2450 mg dm-3 NaCl, pH 7; O 3000 mg dm-3 Na2SO4, pH 7.
electrolyte, aqueous solutions with chloride or sulfate ions were tested. Prior to each experiment, the electrodes were treated with a solution of 1.30 M HCl in order to reject any effect due to the different prehistory of the electrodes. Chemical dissolution experiments were carried out in a batch operation using stirred beakers. These beakers were initially filled with the solution of chloride (or sulfate), and a piece of aluminum was placed inside. Samples were taken from the beakers, and the pH and the aluminum were measured according to the methods previously described. The aluminum dissolution rates were calculated after fitting the experimental data obtained in these essays to the mass balance equations of this batch reactor. The initial and the final weights of the aluminum piece were also used to confirm the results. Results and Discussion Chemical Dissolution of Aluminum. Figures 3 and 4 show, for four different pHs, the aluminum dissolved versus time in a batch reaction system when a sheet of 4 cm2 area (and 0.8 mm thickness) is placed into a solution that contains sodium chloride or sodium sulfate. These figures also show the pH variation with time in each experiment. It can be observed that, in all cases, a chemical dissolution of the aluminum is obtained, with a linear trend for all the experiments except for that corresponding to pH 13, in which the concentration of aluminum reached a final value due to the complete dissolution of the aluminum sheet. The pH decreases in every case, although the change is more significative for a neutral pH. The supporting media seems not to have an important effect on the dissolution results. Figure 5 shows the aluminum dissolution rates as a function of the pH. It can be observed that there is the presence of a minimum at a close to neutral pH and that the dissolution rate is several orders of magnitude higher at alkaline pHs. During the alkaline pH experiments, hydrogen bubbles evolving from the sheet surface were clearly observed.
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Figure 4. Aluminum dissolved and pH variation during batch chemical dissolution experiments carried out under alkaline conditions: [ 2450 mg dm-3 NaCl, pH 12; 0 2450 mg dm-3 Na2SO4, pH 12; 2 2450 mg dm-3 NaCl, pH 13; O 2450 mg dm-3 Na2SO4, pH 13.
Figure 5. Aluminum dissolution rates as a function of the pH. Supporting media: [ 2450 mg dm-3 NaCl; 0 3000 mg dm-3 Na2SO4.
Electrodissolution of the Aluminum Sheets. Figure 6 shows the aluminum, pH, and voltage profiles versus time (dynamic response) generated in typical aluminum electrodissolution experiments (current density ) 0.5 mA cm-2, temperature ) 25 °C, flow rate ) 19.8 dm3 h-1, and pH ) 6.0) carried out in a continuous flow cell. It can be observed that the concentration of aluminum increases with time up to a steady-state value. Likewise, it can be observed that the bulk pH increases and the cell voltage decreases and that both parameters also find a steady-state value. The steady-state value for the concentration of aluminum is greater than the expected value if the process is purely electrochemical. Thus, for the conditions used in the experiments shown in Figure 6, only 0.95 mg dm-3 of electrodissolved aluminum was expected. The pH increase can be justified by the hydroxyl ion generation on the cathode as a consequence of the hydrogen evolution process. The cell voltage decrease can be justified by the increase in the ionic conductivity of the treated solution. Figure 7 shows the variation of the steady-state aluminum concentration electrogenerated in the electrochemical processes with the passed electrical charge,
Figure 6. Concentration of aluminum, pH, and voltage profiles with time (dynamic response) generated in typical aluminum electrodissolution experiments: current density, 0.5 mA cm-2; temperature, 25 °C; flowrate, 19.8 dm3 h-1; pH, 6. Supporting media: [ 2450 mg dm-3 NaCl; 0 3000 mg dm-3 Na2SO4.
Figure 7. Variation of the steady-state aluminum concentration electrogenerated in the electrochemical processes with the electrical charge passed: T, 25 °C; pH, 6.0; flowrate, 19.8 dm3 h-1. Supporting media: [ 2450 mg dm-3 NaCl, 0 3000 mg dm-3 Na2SO4. s Values predicted using Faraday’s Law.
compared with the expected values if the process were purely electrochemical. As can be seen, the electrodissolved aluminum increases linearly with the electrical charge and the experimental values are greater than the values calculated if the process is purely electrochemical. In this figure, it can also be observed that the supporting media does not seem to play an important role in the dissolution process. Figure 8 shows the influence of pH in the steady-state concentration of aluminum. This figure also shows the theoretical value of the aluminum concentration if the process was uniquely electrochemical. It can be seen that the experimental values are greater than the theoretical ones and that the differences increase in highly alkaline media. The discontinuous line corresponds to the amount of aluminum chemically dissolved at the pH of the bulk solution in the cell. It can be observed that, only at pH 12, this line fits with the experimental points. Figure 9 shows the influence of the salinity and that of the current density for the same passed electrical charge. It can be observed that, inside the conductivity range used, the electrodissolved aluminum concentra-
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Figure 8. Influence of pH on the steady-state concentration of aluminum generated in the electrochemical processes: s Faraday’s Law; - - - aluminum chemically dissolved at the pH of the bulk solution in the cell. T, 25 °C; j, 0.5 mA cm-2; flowrate, 19.8 dm3 h-1. Supporting media: [ 2450 mg dm-3 NaCl; 0 3000 mg dm-3 Na2SO4. Figure 10. Main stages involved in the electrochemically assisted coagulation process according to the results obtained in this work.
the cathode, where the pH can become strongly alkaline. This can justify the important contribution of the chemical dissolution to the total dissolution rate (Figure 10). To confirm this role of the chemical dissolution, a model which considers both chemical and electrochemical processes is proposed in the next section and applied successfully to the description of the process. Model of the Electrodissolution Process. From the experimental results described in the former sections, it can be concluded that, to model the electrodissolution process, both chemical and electrochemical processes have to be considered. According to the literature,17 the chemical dissolution process corresponds to the oxidation of the aluminum sheets with the simultaneous reduction of water to form hydrogen, according to eq 1.
2Al + 6H2O f 2Al3+ + 3H2 + 6OHFigure 9. (a) Influence of the salinity on the amount of aluminum electrodissolved: T, 25 °C; pH, 6.0; flowrate, 19.8 dm3 h-1; j, 0.5 mA cm-2. Supporting media: [ NaCl, 0 Na2SO4. (b) Influence of the current density on the amount of aluminum electrodissolved for the same passed electrical charge: T, 25 °C; pH, 6; passed electrical charge, 0.0028 A h dm-3. Supporting media: [ 2450 mg dm-3 NaCl; 0 3000 mg dm-3 Na2SO4.
tion is maintained almost constant. On the other hand, high current densities lead to less efficient processes. This fact can be explained in terms of the competition between aluminum dissolution and oxygen evolution. At low values of current density, only the aluminum dissolution exists, while at high values, both processes compete and, as a consequence, the aluminum generation efficiency decreases. The differences observed between the experimental results and those expected if the process was purely electrochemical can be explained in terms of the chemical dissolution of the electrode surfaces. In the first section of this discussion, it has been proven that chemical dissolution of aluminum can be obtained under acidic or alkaline media, although the rate increases strongly in the later conditions. The electrochemical oxidation and reduction of water can importantly modify the pH on the anode and on the cathode surfaces with respect to the bulk pH. This is specially important on
(1)
On the other hand, the electrochemical processes that occur on the anode and on the cathode surfaces are represented by eqs 2-4. On the anode, aluminum dissolution and oxygen evolution can compete. On the cathode, hydrogen evolution is the main expected reaction. As can be observed, in both the chemical and the
Al f Al3+ + 3e-
(2)
2H2O f O2 + 4H+ + 4e-
(3)
H2O + e- f 1/2H2 + OH-
(4)
electrochemical processes, hydrogen is generated. Once the aluminum is dissolved (chemically or electrochemically), some species can be formed, depending on the pH of the solution. The reactions involved are shown below19,20 (eqs 5-9). In this set of equations, the following have also been taken into account: the carbonate/bicarbonate equilibria (eqs 10 and 11) and the ionization of water, because of its high influence on the calculation of the pH value.
Al(OH)4- + H+ a Al(OH)3
(5)
Al(OH)3 + H+ a Al(OH)2+
(6)
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Al(OH)2+ + H+ a Al(OH)+2
(7)
Al(OH)+2 + H+ a Al3+
(8)
Al(OH)3(s) a Al3+ + 3OH-
(9)
CO2 + H2O a H+ + HCO3-
(10)
HCO3- a CO32- + H+
(11)
H2O a OH- + H+
(12)
The other important point in the model is the fluid dynamic approach. As has been shown previously, the pH has a strong influence in the chemical dissolution process, and the value of this parameter can increase by several order of magnitude at an alkaline pH. In each electrochemical cell, there is a pH profile between anode and cathode. In the anode, the water oxidation process (oxygen evolution) generates a high concentration of protons and, thus, a lower pH must be obtained. In the cathode, the water reduction process results in the formation of hydroxyl ions and a higher pH must appear in this zone. Thus, a marked pH profile is expected, and the bulk pH is not a good value to determine the rate of chemical dissolution, because it must differ considerably from the actual values on the electrodes surfaces. To take this into account, the fluid dynamic model must consider this profile of pH. An exhaustive description of the electrochemical dissolution of aluminum electrodes (in which the concentration profile of every compound in the electrochemical cell is calculated) is extremely difficult, because it would lead to a mathematical system involving several partial differential equations. This complex situation arises since the concentration of every compound depends on time and on the distance to the electrode surface. Such a system would not be helpful because of its high degree of complexity, and therefore, in an attempt to achieve a useful model, a number of assumptions must be made in order to simplify the model. The assumptions that are going to be taken into account have been previously proposed in the literature in other works of our group,21,22 and they have helped to successfully represent wastewater electrochemical oxidation processes.23 The first assumption is related to the description level of the process. The position-dependence of the model can be simplified by dividing the electrochemical reactor into three zones: two zones (electrochemical zones) close to the electrodes (anode and cathode) and a third zone corresponding to the bulk solution (chemical zone). In these three zones, the concentration of every compound is considered to be constant with position and only timedependent. This assumption is valid if the residence time in the electrochemical cell is small, since, in this case, the profiles of concentration in the flow direction can be assumed to be negligible. The concentration of each compound in the chemical zone is taken as the value measured experimentally. The concentration in the electrochemical zone is supposed to have a value between the concentration at the electrode surface and the concentration in the chemical zone. Hence, the electrochemical reactor is modeled as the combination of three interconnected stirred-tank reactors, as shown schematically in Figure 11. This highly
Figure 11. Scheme of the electrochemical reactor considered in the model as the combination of three interconnected stirred-tank reactors.
segregated flow approach has been successfully demonstrated in the literature using marker-pulse techniques for channel flow cells24. The volume of each zone can be easily calculated25,26 if it is assumed that the thickness of the electrochemical zone is equivalent to the Nernst diffusion layer (δ). This assumption is acceptable because direct oxidation and most of the mediated oxidation processes (those with high reaction rates) occur in this zone. The thickness of this zone can be evaluated as a function of the masstransfer coefficient (k) and the diffusivity (D) using eq 13.
δ)
D k
(13)
The electrode surface area is known, and so the volume of each electrochemical zone (Va and Vc) can be easily calculated by multiplying it by the thickness (δ). The remaining volume of the system corresponds to the volume of the chemical zone (Vb). Mass transport processes between electrochemical and chemical zones are quantified by assuming that the local exchange rate is proportional to the concentration difference between the two zones. The mass-transfer rate can be calculated from eq 14, where [Si]* and [Si] are the concentrations (mol m-3) of component i in the electrochemical and the chemical zones, respectively, k (m s-1) is the mass-transfer coefficient, and A (m2) is the specific interfacial area between the electrochemical and chemical zones. The mass-transfer coefficient (k) can be assumed to depend only on the flow rate conditions, because the concentrations of the compounds are low.
ri ) kA([Si]* - [Si])
(14)
The assumption of dividing the electrochemical cell into several zones allows the simplification of the mathematical complexity of the model. Thus, the complex system of partial differential equations (obtained from the mass balance in a nonsimplified system) is reduced to an easier-to-solve system of ordinary differential equations.
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Va
d[Si]a ) kA([Si]b - [Si]a) + relectrochemi,a + rchemi,aVa dt (15)
[Al(OH)3]z[H+]z
Vc
d[Si]c ) kA([Si]b - [Si]c) + relectrochemi,c + rchemi,cVc dt (16)
[Al(OH)2+]z[H+]z
d[Si]b ) q([Si]0 - [Si]b) + kA([Si]a - [Si]b) + dt kA([Si]c - [Si]b) + rchemi,bVb (17)
[Al(OH)2+]z[H+]z
Vb
[Al(OH)2+]z
[Al(OH)2+]z
3+
[Al ]z [HCO3-]z[H+]z
The kinetic rate of the electrochemical processes, limited by the applied current density (I), is given in eq 18, where F is the Faraday constant.
I r) F
I ηi F
∑j ηj
(19)
To obtain the chemical dissolution rate that appears on both equations, the data of Figure 5 have been fitted to an empirical equation. The Euler method was used to solve the system of differential equations. The aluminum dissolved from the electrodes can be present as different species depending on the pH of the media. The concentration of any of these species depends on the total concentration of dissolved aluminum and on the pH, and this makes the system complex from the mathematical point of view and, consequently, difficult to solve. To simplify the calculations, eqs 15-17 were applied only to a unique aluminum species (the total dissolved aluminum, TDA) and to the hydroxyl and protons. For each time step (of the Euler differential-equations-solving method), the different aluminum species and the resulting proton and hydroxyl concentration in each zone were recalculated using a pseudoequilibrium approach. To do this, the equilibrium equations (eqs 20-27), the charge (eq 28), and the aluminum (eq 29) and inorganic carbon (IC) balances (eq 30) were considered in each zone (anodic, cathodic, and chemical), and a nonlinear iterative procedure (based on an optimization method) was applied to satisfy simultaneously all the equilibrium constants. In these equations (eqs 20-30), subindex z stands for the three zones in which the electrochemical reactor is divided (anodic, cathodic, and chemical).
[Al(OH)4-]z[H+]z [Al(OH)3]z
[CO32-]z[H+]z -
(18)
On the cathode, only hydrogen evolution is supposed to develop. Thus, eq 18 can be used to represent this process. On the anode, both aluminum electrodissolution and oxygen evolution compete. To model this, the fraction of current employed to oxidize each of these compounds can be assumed to be directly related to the overpotential of the two compounds (ηAl and ηO2). This approach has been used successfully as reported in the literature in previous works of our group.21-23
relectrochemi )
[H2CO3]z
) 10-8.0
[Al3+]z[OH-]z3 ) 10-32.9
(20) (24)
[HCO3 ]z
) 10-5.7
(21)
) 10-4.3
(22)
) 10-5.0
(23)
) 10-6.37
(25)
) 10-10.25
(26)
[H+]z[OH-]z ) 10-14.0 n[An-]z m[Bm+]z ) 0 ∑ ∑ anionic species cationic species
(27) (28)
[TDA]z ) [Al3+]z + [AlOH2+]z + [Al(OH)2+]z + [Al(OH)3]z + [Al(OH)3]sol,z + [Al(OH)4-]z (29) [IC]z ) [H2CO3]z + [HCO3-]z + [CO32-]z
(30)
Figure 12 shows the time variation of pH and aluminum concentration calculated from the model compared to those obtained experimentally for three different experimental assays. It can be observed that the model fits well the experimental data, with average errors
