Anodic Dissolution of Pure Aluminum during Electrocoagulation

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Anodic Dissolution of Pure Aluminum during Electrocoagulation Process: Influence of Supporting Electrolyte, Initial pH, and Current Density Khaled Mansouri,† Karim Ibrik,‡ Nasr Bensalah,*,†,‡ and Ahmed Abdel-Wahab‡ † ‡

Department of Chemistry, Faculty of Sciences of Gabes, University of Gabes, Cite Erriadh, Zrig 6072, Gabes, Tunisia Department of Chemical Engineering, Texas A&M University at Qatar, Education City, P.O. Box 23874, Doha, Qatar ABSTRACT: In this work, effects of some experimental parameters (supporting electrolyte, initial pH, and current density) on aluminum corrosion and anodic dissolution of aluminum during electrocoagulation process were investigated. Potentiodynamic polarization tests, impedance spectroscopy measurements and potentisotatic current density transients were used to evaluate corrosion parameters and verify effects of supporting electrolyte and initial pH on aluminum corrosion. The presence of NaCl and Na2SO4 shifted the corrosion potential toward more cathodic potentials, indicating that corrosion of aluminum is catalyzed by the presence NaCl and Na2SO4. On the contrary, the presence of NaH2PO4 increased the corrosion potential, which indicates that the presence of NaH2PO4 inhibits the corrosion of aluminum. Galvanostatic electrolyses demonstrated that measured concentrations of aluminum exceeded theoretical values calculated using Faraday’s Law. The excess in dissolved aluminum produced during galvanostatic electrolyses is primary due to the chemical dissolution of aluminum, which is more significant at highly alkaline conditions. A final pH of value around pH 9 was observed in the presence of NaCl for pH values in the range pH 411 which can be explained by buffering effects of aluminum hydroxides.

1. INTRODUCTION The passivity of aluminum sacrificial electrodes was extremely problematic for the treatment of industrial effluents by electrocoagulation.14 Passive films are readily formed on the surface of aluminum and its alloys in open air and in water.5,6 The corrosion resistance of aluminum depends largely on the formation of a native oxide film on its surface. The aluminum oxide film is relatively chemically inert, exhibiting the passive behavior of aluminum. The formation of this passive film that adheres to the anode surface prevents the dissolution of aluminum and restricts charge transfer between solution and electrodes. Passivation of sacrificial anodes leads to excessive consumption of electricity and reduces the efficiency of wastewater treatment by electrocoagulation. The rate of corrosion of sacrificial aluminum anodes depends mainly on two mechanisms: formation and build-up of a passive aluminum-oxide layer and subsequent partial destruction of that layer through pitting.712 Aluminum is much more susceptible to pitting than other metals, and a variety of anions that may exist in the solution can bring about pitting. In most practical cases of aluminum corrosion, the chemical reaction of anions adsorbed on the surface of the aluminum oxide film with cationic Al species in the oxide lattice significantly affect the overall aluminum dissolution reaction. This chemical reaction causes the local breakdown of the passive oxide film, which is often followed by pitting corrosion.1113 Few researchers1421 have suggested the addition of supporting electrolytes, such as sodium chloride (NaCl), to reduce the passivation of working electrodes and to increase the conductivity of the medium. Additionally, application of a sinusoidal intensity to prevent the formation of a deposit on electrode r 2011 American Chemical Society

surface has also been suggested by some researchers.1416 However, literature lacks attention to the effect of experimental parameters such as nature and concentration of supporting electrolyte, pH, and current density on the rate of destruction of the aluminum protective film that forms on sacrificial aluminum anodes, which in turn influences the overall rate of the electrochemical reactions and mechanisms involved in electrocoagulation.2227 The aim of this study is to obtain information on the corrosion behavior of aluminum sacrificial anodes and evaluate the effects of supporting electrolyte nature and concentration, pH, and current density on electrochemical dissolution of aluminum. Two main experimental sets were carried out to achieve the objective of this study. First, the corrosion behavior of an aluminum anode in different electrolytes (NaCl, Na2SO4, and NaH2PO4) was examined using potentiodynamic polarization tests. Second, the effect of pH and current density parameters on the electrochemical dissolution of aluminum was studied using galvanostatic electrolysis of aluminum electrodes.

2. MATERIALS AND METHODS 2.1. Chemicals and Analytical Methods. All reagents used were of analytical grade obtained from Fulka or Acros. Relevant aqueous solutions were prepared by dissolving lab grade reagents (NaCl, Na2SO4, NaH2PO4) in deionized water obtained from a Received: June 6, 2011 Accepted: October 5, 2011 Revised: October 3, 2011 Published: October 05, 2011 13362

dx.doi.org/10.1021/ie201206d | Ind. Eng. Chem. Res. 2011, 50, 13362–13372

Industrial & Engineering Chemistry Research Milli-Q system (Molsheim, France), with resistivity >18 mΩ at 25 °C. Working electrodes were comprised of aluminum plates cut from a pure aluminum sheet (99.99% purity) of 2 mm thickness. Prior to each experiment, the aluminum working electrode was manually finished using successively finer grades of abrasive paper, polished with 0.3 μm alumina, rinsed with 1.5 M HCl followed by rinsing with deionized water, and then dried before immersion in the electrolyte solution. Samples were taken periodically and analyzed for dissolved aluminum concentration and pH value. Concentration of dissolved aluminum during galvanostatic electrolysis was determined by atomic absorption spectrometry (Zeeman spectrophotometer HITACHI Z-6100) after acidification using nitric acid (4 N HNO3 ). The pH of the aqueous solutions was measured using inoLab WTW pH-meter, equipped with a combined glass electrode (METROHM). The conductivity was measured using a conductivity meter (MeterLab, type CDM230) equipped with a conductivity cell comprised of two platinum electrodes. Morphological observations of the aluminum electrode surfaces were conducted using a metallurgical trinocular microscope (AuxiLab, model ZUZI 173/2) after galvanostatic electrolyses. 2.2. Potentiodynamic Polarization and Potentiostatic Current/Time Measurements. Potentiodynamic tests were used to evaluate the pitting behavior and passivity of aluminum in 0.10.4 M NaCl, 0.1 M Na2SO4, and 0.1 M NaH2PO4 aqueous solutions. Potentiodynamic polarization experiments of aluminum were performed with a three-electrode cell (Radiometer C145/170) using a potentiostat/galvanostat (PGZ 301 20 V 1A) controlled by Voltalab 40 software to allow data acquisition. A platinum plate and a saturated calomel electrode (SCE) were respectively used as auxiliary and reference electrodes. The working electrode was a 1 cm2 pure aluminum plate (99.99%) inserted into a PTFE sample holder (Radiometer PEK 29). The electrochemical potential was scanned positively starting from 2.0 V up to +2.0 V vs SCE at a scan rate of 0.5 mV s1. The potentiostatic current/time transient measurements were carried out after a two step procedure as follows: the working electrode was first held at the starting potential 2.0 V/SCE for 1 min to attain a reproducible electroreduced electrode surface; then the electrode was suddenly polarized in the positive direction to a potential equal Epit for 30 min, after that current transient was recorded. 2.3. Electrochemical Impedance Spectroscopy. Electrochemical impedance spectroscopy (EIS) measurements were carried out using a potentiostat/galvanostat (PGZ 301 20 V 1A) controlled by Voltalab 40 software to allow data acquisition of corrosion potentials (Ecorr) of aluminum electrodes immersed in aerated solutions. After determination of steady-state current at Ecorr, peak-to-peak sine wave voltages (10 mV) were superimposed on the resting potential at frequencies between 100 kHz and 10 mHz. The EIS diagrams are shown in Nyquist form (shown where. If they are shown in a figure, please indicate the figure no.). 2.4. Galvanostatic Electrolyses. Anodic dissolution of aluminum was carried out under galvanostatic conditions. The electrolyses experiments were performed in a one compartment thermostatted electrolytic cell (parallel rectangular aluminum electrodes, 24 cm2, interelectrode gap 2 cm). Aluminum electrodes were connected to a digital DC power supply (Monacor PS430) operating under galvanostatic options, providing current and voltage in the range of 030 A and 020 V, respectively. The cell voltage was recorded using a potentiometer. During the

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Figure 1. Changes of open circuit potential with time during immersion of aluminum plate in aerated deionized water. Operating conditions: anode: pure Al specimen (1 cm2); reference electrode: SCE; T = 25 °C.

galvanostatic electrolyses process, the aqueous solution (0.2 dm3) was vigorously stirred using a magnetic stir bar (200 rpm), and the temperature was held constant at 25 °C. Aqueous solution pH was adjusted using 0.1 M NaOH or HCl to achieve the desired initial pH values for each experiment.

3. RESULTS AND DISCUSSION 3.1. Study of Aluminum Corrosion by Potentiodynamic Polarization. In order to better understand the influence of

supporting electrolyte on the electrochemical corrosion of aluminum, open circuit potential and potentiodynamic polarization experiments were performed in an attempt to explain the phenomena occurring on the anode surface. Figure 1 displays the changes of an aluminum electrode in the open-circuit potential (OCP) with immersion time in deionized water at room temperature. The OCP instantly increased from an initial value of 795 mV/SCE to a value of 670 mV/SCE, and it remained constant at this value with minor fluctuations with time around this value. The immediate increase of OCP can be explained by the growth of Al(OH)3 film on the surface of the Al electrode. The fluctuation of the OCP can be attributed to the changes of pH near the Al electrode surface. Based on these results, a corrosion potential value (Ecorr) of 670 ( 5 mV/SCE can be deduced for pure Al in aerated aqueous solutions. This value is very similar to those reported in the literature.28,29 Figure 2 shows anodic and cathotic potentiodynamic polarization curves obtained from pure Al plates at a scan rate of 0.5 mV s1 in aerated aqueous solution without supporting electrolyte. As can be seen from Figure 2, both anodic and cathodic branches display a typical Tafel behavior. Accurate evaluation of corrosion potential (Ecorr) and corrosion current density (jcorr) was made using the Tafel extrapolation method30,31 and presented in Figure 2. The anodic polarization curve shows an extremely active dissolution followed by a limitation at high current densities. Corrosion potential (Ecorr) of aluminum as determined by the Tafel extrapolation method is very comparable to that determined previously by open circuit potential measurements: Ecorr= 668 mV/SCE. Corrosion current (jcorr) of aluminum electrode was equal to 2.6 μA cm2. Figure 3-a shows potentiodynamic polarization curves obtained from an aluminum plate in aerated aqueous solutions containing 0.1 M of NaCl, Na2SO4, or NaH2PO4 under alkaline conditions (pH = 9) at a scan rate of 0.5 mV s1. The potentiodynamic 13363

dx.doi.org/10.1021/ie201206d |Ind. Eng. Chem. Res. 2011, 50, 13362–13372

Industrial & Engineering Chemistry Research

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Figure 2. Potentiodynamic polarization curve of aluminum in aerated deionized water, and extrapolation Tafel method. Operating conditions: anode: pure Al specimen (1 cm2); cathode: Pt wire; reference electrode: SCE; pH = 7; scan rate: 0.5 mV s1; T = 25 °C.

polarization plot without supporting electrolyte was also presented in the same figure for comparison (in this case the current density is expressed in μA cm2). Figure 3 shows that the anodic electrode behavior in the presence of supporting electrolyte was completely different from that observed in aerated deionized water without supporting electrolyte. Also, the behavior of the anodic polarization branches in the presence of supporting electrolyte do not display the expected log/linear Tafel behavior over the whole applied range of potential. The observed curvature of the anodic branch may be attributed to the deposition of corrosion products on the aluminum surface to form a nonpassive surface film. Aluminum showed a classical passive region in which current remained almost unaffected by the change in applied potential. This can be observed as a current density plateau in Figure 3-b. However, the current density increased abruptly after reaching a certain value of electrode potential which is the pitting potential (Epit). The pitting corrosion was preceded by uniform thinning of the hydroxide/oxide protective film that prevails over pitting corrosion prior to the pitting potential. After reaching Epit, current density continued to increase very slightly with increasing potential. As can be observed, the nature of the electrolyte greatly affects corrosion current density and corrosion potential. The values of corrosion current density and corrosion potential in solutions containing different electrolytes are shown in Table 1. The presence of NaCl or Na2SO4 shifted the corrosion potential toward more negative values, indicating that corrosion of aluminum is catalyzed by the presence NaCl or Na2SO4. In contrast, the presence of NaH2PO4 increased the corrosion potential, which indicates that the presence of NaH2PO4 inhibits the corrosion of aluminum. Moreover, the experimental results indicated that the corrosion rate of aluminum is highest in the presence of NaCl because the corrosion current density is the largest. Figure 4 shows the Nyquist plot obtained on pure aluminum in aqueous solutions containing 0.1 M of NaCl, Na2SO4, or NaH2PO4. The shapes of the Nyquist plots corresponding to these conditions include one small semicircle at the high frequency region, one depressed semicircle, and a diffusion tail which appears inclined to the Real Impedance axis at 45° at the low frequency region. In general, the capacitance no longer affects the total impedance value in the diffusion tail portion. The diffusion process can be

Figure 3. (a) Potentiodynamic polarization curves of aluminum in aerated aqueous solutions containing different supporting electrolytes. (b) Different parts of polarization curve: charge transfer, anodic passivation, and pitting corrosion. Operating conditions: anode: pure Al specimen (1 cm2); cathode: Pt wire; reference electrode: SCE; pH = 9; electrolyte concentration: 0.1 M; scan rate: 0.5 mV s1; T = 25 °C.

caused by the diffusion of Al3+ ions, which is produced by the anodic reaction, from the metal surface to the bulk solution. Therefore, the results presented in Figure 4 show that the diffusion tail at impedance plot which indicates the beginning of metastable pitting. In addition, the depressed semicircle could be due to microscopic surface roughness and the presence of a porous corrosion product film. The diameter of low frequency semicircle changes among supporting electrolytes. The diameter of the semicircle in the presence of NaCl is smaller than that in the presence of NaH2PO4 or Na2SO4 which suggests that the passive film formed on pure aluminum in the presence of NaH2PO4 or Na2SO4 is more protective to corrosion. This is in accordance with the results of polarization curves. These observations can be summarized as follows: - The addition of supporting electrolyte increases the ionic conductivity of the solution, which is a determining factor in any electrochemical process. The increase in conductivity has a significant influence on the kinetic parameters of electrochemical reactions and it increases the corrosion rate of aluminum. - For alkaline pH conditions, the corrosion of aluminum leads to the formation of a porous layer consisting of amorphous Al(OH)3 and crystalline Al2O3 on the surface of the anode.4,17,26 This layer is characterized by a heterogeneous structure and irregular pores. This layer significantly reduces the corrosion rate of aluminum, but it cannot fully protect the aluminum surface. The increase of anode potential damages this layer 13364



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Table 1. Results of Potentiodynamic Polarization of Aluminum Determined by the Tafel Extrapolation Method in the Presence of Different Supporting Electrolytesa electrolyte

corrosion potential

corrosion current

Ecorr (mV)

density jcorr (μA cm2)

without supporting electrolyte

668

2.6

0.1 M NaCl

1212.7

21

0.1 M Na2SO4

1056.0

8.7

0.1 M NaH2PO4

518.5

4

a

Experimental conditions: pH = 9; electrolyte concentration: 0.1 M; scan rate: 0.5 mV s‑1; T = 25 °C.

Figure 4. Electrochemical impedance spectroscopy spectra obtained on pure aluminum in aerated aqueous solutions containing different supporting electrolytes. Operating conditions: anode: Al specimen (1 cm2); cathode: Pt wire; reference electrode: SCE; pH = 9; electrolyte concentration: 0.1 M; scan rate: 0.5 mV s1; T = 25 °C.

through the diffusion of anionic species toward the anode by migration. Anionic exchange reactions between hydroxide ions and anions of the supporting electrolyte occur and form soluble aluminum complexes. This exchange also causes the direct dissolution of the passive layer of Al(OH)3 and Al2O3. All of these phenomena constitute the pitting corrosion of aluminum. Pitting corrosion is responsible for the increase in current density right after the passivation plateau, as shown in Figure 3-b. The pitting potential of aluminum is the potential of passivation breakdown for which there is an acceleration of the corrosion process by diffusion of anions through the pores of the Al(OH)3 layer. It appears that pitting corrosion depends greatly on the nature of supporting electrolyte because the value of potential Epit varies among different electrolytes. In the presence of NaCl, pitting corrosion occurs at a much lower potential than that in the presence of Na2SO4 or NaH2PO4 electrolytes. This is most likely due to the size of chloride ions, which are relatively smaller than sulfate and phosphate ions. Figure 5-a and -b provides typical Nyquist plots and potentiostatic anodic current transients, respectively obtained from pure aluminum specimen at a potentials equal to open circuit potential in aerated aqueous solutions containing different NaCl concentrations. The results presented in Figure 5-a show that the diameter of low frequency semicircle increased with NaCl concentration suggesting that the passivity of the oxide film decreases with NaCl concentration increase. However, the anodic current density shown

Figure 5. (a) Electrochemical impedance spectroscopy spectra and (b) potentiostatic current density transients at pitting potentials obtained from pure aluminum in aerated aqueous solutions containing different concentrations of NaCl. Operating conditions: anode: pure Al specimen (1 cm2); cathode: Pt wire; reference electrode: SCE; pH = 9; scan rate: 0.5 mV s1; T = 25 °C.

in Figure 5-b initially remained constant with time at a value close to zero up to an induction time of about 600 s, and then it rapidly increased to a certain value at which it remained constant. The ascending current after induction time can be attributed to the growth of pit. It is noted that the pitting of pure aluminum increased by increasing chloride ions concentration. These ions are characterized by their ability to penetrate through the Al(OH)3/Al2O3 film forming soluble complexes.2628 The increase of chloride concentration further decreases corrosion potential and minimizes the difference between pitting potential and corrosion potential as can be observed from Table 2. At a potential close to Epit, Cl ions are strongly adsorbed, and polarization curves show an increase in the current density. This indicates an enhanced promotion of the passive film hydration and hence its degradation as described by the following reactions Al2 O3 þ 6Cl þ 6Hþ / 2AlCl3 þ 3H2 O AlðOHÞ3 þ Cl / AlðOHÞ2 Cl þ OH AlðOHÞ2 Cl þ Cl / AlðOHÞCl2 þ OH AlðOHÞCl2 þ Cl / AlCl3 þ OH AlCl3 þ Cl / AlCl4 13365



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Table 2. Corrosion and Pitting Potentials Obtained from Potentiodynamic Polarization Curves of Aluminum in the Presence of NaCl at Different Concentrationsa corrosion potential pitting potential

a

electrolyte

Ecorr (mV)

Epit (mV)

Epit-Ecorr (mV)

without electrolyte

668.0

-

-

0.1 M NaCl

1212.7

646.0

566.7

0.15 M NaCl

1223.8

662.0

561.8

0.2 M NaCl

1255.5

707.0

548.5

Experimental conditions: pH = 9; scan rate: 0.5 mV s1; T = 25 °C.

In the presence of sulfate ions, the observed pitting was much lower than that observed in the presence of chloride ions. In the case of solutions containing sulfate, the passive layer of Al(OH)3/ Al2O3 is partially dissolved, and consequently the rate of electrochemical corrosion of aluminum is lower. Whereas in the presence of phosphate ions, no pitting was observed due to the formation of insoluble complexes of AlPO4 that further inhibit the corrosion of aluminum.3235 Figure 6-a shows potentiodynamic polarization curves for pure aluminum immersed in 0.1 M NaCl aerated aqueous solutions at different pH values of 2, 5, 7, 9, and 11. Similar behavior of potentiodynamic polarization was observed for all tested pH values. However, a shift of corrosion potential was observed when pH was changed from neutral pH to acidic or alkaline pH values. Indeed the corrosion potential in neutral solution is the lowest. The beginning of pitting is not visible in these cases since the pitting potentials Epit coincided with the corrosion potentials Ecorr. From these results, it can be concluded that the main degradation mechanism in 0.1 M NaCl aerated aqueous solutions of pH values 2, 5, and 7 is pitting corrosion and the pitting potential is independent of pH over this pH range. On the contrary, in case of alkaline solutions (pH values of 9 and 11), the values of corrosion potential Ecorr were observed to be lowered to more cathodic values without a change in the pitting potential. Further, the decrease in the corrosion potential Ecorr values, compared to those of acidic and neutral pH, indicate the loss of passivity of aluminum due either to thinning of surface oxide layer by hydroxide ions (OH) attack (alkaline chemical dissolution) or to the absence of the primary oxide film.17,26 EIS spectra of Figure 6-b show that increasing pH from 2 to 7 has no influence on Nyquist plots. However a decrease of diameter of semicircle at high frequency was observed for pH 9 and 11. This confirms that thinning of surface oxide layer by hydroxide ions occurs at alkaline conditions. Furthermore, morphological observations by optical microscope of aluminum electrode surfaces before and after anodic polarization at a constant current density of 10 mA cm2 in aqueous solutions containing 0.1 M of NaCl, Na2SO4, or NaH2PO4 at pH 9 are shown in Figure 7. This figure shows that large localized pitting corrosion was clearly observable on the surface of aluminum plate polarized in the presence of NaCl. However, in the case of Na2SO4, uniform distributed pitting corrosion was observed. In contrast, the aluminum plate polarized in the presence of NaH2PO4 was completely covered with a white porous layer, which is most likely an insoluble layer of AlPO4. It appears that this layer was of different nature and structure and was thicker than the layers observed in the presence of the other two electrolytes.

Figure 6. (a) Potentiodynamic polarization curves and (b) electrochemical impedance spectroscopy spectra obtained on pure aluminum in NaCl aerated aqueous solutions at different initial pH values. Operating conditions: anode: pure Al specimen (1 cm2); cathode: Pt wire; reference electrode: SCE; electrolyte concentration: 0.1 M; scan rate: 0.5 mV s1; T = 25 °C.

These results indicate the significance of supporting electrolyte in aluminum corrosion. It has been observed that the corrosion of aluminum is more favorable in the presence of chloride ions than that in the presence of sulfate or phosphate ions. The phosphate ions inhibit corrosion of aluminum through the formation of a protective layer of aluminum phosphate (phosphatation). 3.2. Electrochemical and Chemical Dissolution of Aluminum. According to the literature,3638 efficiency of the electrocoagulation process in treating wastewaters depends largely on the amount and speciation of aluminum dissolved during the electrochemical treatment process. In order to evaluate the dissolution of aluminum during galvanostatic electrolyses, the amount of aluminum dissolved (mexp) was measured and compared to the calculated theoretical amount expected to be dissolved (mth) using Faraday’s law. Figure 8 presents comparison between theoretical and experimental aluminum concentration dissolved at different specific electrical charges (A.h.L1) during galvanostatic electrolysis of 0.1 M NaCl aerated aqueous solution using aluminum electrodes at a current density of 10 mA cm2. Theoretical aluminum concentration (in g L1) was calculated using Faraday’s law, assuming that the provided electrical energy was entirely used only to oxidize Al to Al3+ according to the following reaction: Al / Al3+ + 3e. 13366



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Figure 7. Optical microscope photos (400) of aluminum morphology before and after anodic polarization at current density j = 10 mA cm2 in aqueous solutions containing 0.1 M NaCl, Na2SO4, and NaH2PO4 at pH 9.

According to Faraday’s law, the electrical charge “q” is calculated using the following equation q ¼ 3F

mAl q MAl w ½Al w mAl ¼ MAl 3F

q MAl 3FV During galvanostatic electrolyses, the current density is constant; thus electrical charge (q) and specific electrical charge (Q) are defined respectively by the following equations: q(A.s) = I t and Q(A.h.L1) = (I t)/(3600 V). The theoretical amount of dissolved aluminum is then given by the following equation ¼

½Alðg L-1 Þ ¼

I t MAl 3600 Q MAl ¼ 3FV 3F

¼ 0:366 Q ðA h L1 Þ where [Al] is the amount of aluminum generated (in g dm3); I is the current intensity (in A); t is electrolysis-time (in s); M is the molecular weight of aluminum (M = 26.98 g mol1); F is the Faraday constant (F = 96487 C mol1); V is the electrolyte volume (V = 0.2 dm3); and Q is the specific electrical charge (in A h dm3).

Figure 8. Changes of theoretical and experimental concentrations of aluminum dissolved during galvanostatic electrolysis with specific electrical charge. Operating conditions: initial pH = 7, j = 10 mA cm2, supporting electrolytes: 0.1 M NaCl.

As can be seen from Figure 8, both theoretical and experimental amounts of aluminum at a current density of 10 m A cm2 increase linearly with increasing specific electrical charge for Q e 2.5 A h L1. At the beginning of galvanostatic electrolysis, 13367



Figure 9. (a) Influence of current density on the changes of aluminum concentration dissolved during galvanostatic electrolyses with specific electrical charge. (b) Photos of samples taken during anodic dissolution of aluminum. Operating conditions: initial pH = 7, j = 10 mA cm2, supporting electrolytes: 0.1 M NaCl.

measured concentrations of aluminum correlate well with theoretical concentrations calculated from the aforementioned equations. However, for specific electrical charge greater than 0.5 A h L1, experimental values exceeded theoretical values and the gap between the two broadened with the increase of specific electrical charge. It should be noted that similar results were reported by other authors.3639 Figure 9-a demonstrates the influence of current density on the change of measured aluminum concentration with specific electrical charge during galvanostatic electrolyses of 0.1 M NaCl aqueous aerated solutions at pH 7. As can be seen, different current densities did not significantly affect the amount of aluminum dissolved by anodic oxidation of the metal. Indeed, comparable amounts of aluminum were measured at different current densities after consuming the same amount of electrical specific charge. This could be explained by the fact that the increase in current density lead to the involvement of secondary reactions that can occur in parallel with anodic dissolution of metal, such as oxygen evolution and formation of hypochlorite ions from anodic oxidation of chloride ions.2527 These side reactions decrease the effectiveness of anodic dissolution of aluminum. Furthermore, it was observed that increasing the current density resulted in an increased rate of hydrogen production and a decrease in size of H2 bubbles generated at the cathode. Additionally, Figure 9-b confirmed the formation of Al(OH)3 solid precipitate during anodic dissolution of aluminum. Generated Al3+ ions react with hydroxide ions (OH) to produce Al(OH)3 solid, which was visually observed in the reactors. The amount of white precipitate increased with time, which is in

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Figure 10. (a) Change of aluminum concentration with time during immersion of aluminum plates (24 cm2) in 0.1 M NaCl aqueous solutions at pH 9. (b) Changes of dissolved aluminum concentration with pH during 60 min immersion of aluminum plates (24 cm2) in 0.1 M NaCl aqueous solution.

agreement with the measured amount of aluminum dissolved during galvanostatic electrolyses. These results proved that galvanostatic electrolysis of 0.1 M NaCl aqueous solutions using aluminum electrodes at neutral pH leads to the formation of Al3+ ions at the anode and the production hydrogen gas at the cathode. The gap between measured aluminum concentrations during galvanostatic electrolysis and the theoretical amounts of aluminum as shown in Figure 8 can be attributed to chemical dissolution of aluminum electrodes by corrosion in addition to anodic dissolution of aluminum. Several researchers have also reported the same behavior.1221 However, to confirm this assumption, it is necessary to evaluate the chemical dissolution of aluminum in water. A set of experiments was conducted in which aluminum plates of known weights were immersed in slightly alkaline 0.1 M NaCl aqueous solutions (pH = 9). These conditions were selected to mimic pH conditions near the cathode where the chemical dissolution of aluminum is most likely to occur. Figure 10 shows changes of aluminum concentration with time, as determined by the gravimetric method during immersion of identical aluminum plates (24 cm2) in 0.1 M NaCl aqueous solutions at pH 9. As can be observed, aluminum concentration increased linearly with time up to 480 min, after which it reached a plateau at about 0.014 g L1 for the rest of the operating time. This behavior may be explained by the formation of a protective layer of aluminum hydroxide/oxide (Al(OH)3/ Al2O3) on the aluminum electrode surface, which inhibits further chemical dissolution of aluminum.14 The formation of this 13368



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protective layer in aerated solutions can be given by the following reactions Al / Al3þ þ 3e 2H2 O þ 2e / H2 þ 2OH The overall chemical reaction is 2Al þ 3H2 O f 2AlðOHÞ3ðsdÞ þ H2 Also, in the presence of dissolved oxygen aluminum can be directly oxidized into aluminum oxide, as shown by the following reaction 4Al þ 3O2ðaqÞ f 2Al2 O3ðsdÞ These results confirmed that aluminum can be chemically dissolved, but the amount dissolved by chemical oxidation with dissolved oxygen at pH 9 is negligible compared to that dissolved by anodic oxidation. As such, chemical dissolution under these operating conditions is not sufficient to explain the gap observed in Figure 8. The gap could be possibly explained by the fact that the operating conditions present during the chemical dissolution of aluminum are different from the actual hydrodynamic conditions and pH in the vicinity of aluminum electrodes used during galvanostatic electrolysis. Particularly, the pH conditions at the cathode are strongly alkaline due to the reduction of water into H2 forming OH ions. To better understand the significance of chemical dissolution of aluminum in the electrocoagulation process, it is necessary to evaluate the influence of initial pH on the chemical dissolution of aluminum. Figure 10-b presents concentrations of dissolved aluminum at various initial pH values during 60 min immersion of identical aluminum plates (24 cm2) in 0.1 M NaCl aerated aqueous solutions. This figure clearly shows that pH had a significant influence on the amount of chemically dissolved aluminum. The chemical dissolution of aluminum was negligible for pH < 10. However, for higher pH values, the concentration of aluminum dissolved in aqueous solutions increased with an increasing pH value from approximately 0.03 g L1 at pH 11 to 0.5 g L1 at pH 12.5. These results demonstrate that the chemical dissolution of aluminum in the electrocoagulation process depends largely on pH in the vicinity of electrodes. Nevertheless, it is difficult to identify exactly the pH conditions at electrodes’ surfaces, which makes it complex to accurately quantify the amount of aluminum chemically dissolved. These results indicate that dissolved concentrations of aluminum measured during galvanostatic electrolyses using aluminum electrodes are attained by both electrochemical and chemical dissolution. The importance of chemical dissolution depends largely on pH conditions close to electrodes’ surfaces. Superfaradic excess of dissolved aluminum is primarily related to the chemical dissolution of aluminum, which is more significant at pH > 11. Canizares et al.3639 also demonstrated that the rate of chemical dissolution of aluminum in water is pH dependent, and it is higher for pH values above 12. 3.3. Influence of Experimental Parameters on Solution pH during Galvanostatic Electrolyses Using Aluminum Electrodes. The preceding results showed that solution pH influences the rate of aluminum dissolution during galvanostatic electrolysis of aqueous solutions using aluminum electrodes. This led us to investigate the influence of experimental parameters on the changes of pH during galvanostatic electrolysis between two aluminum

Figure 11. Influence of (a) initial pH, (b) current density, and (c) supporting electrolyte on changes of pH vs time during galvanostatic electrolysis using aluminum electrodes.

plates. Based on the literature,4045 it was reported that initial pH, current density, and nature of supporting electrolyte are the main parameters that can cause pH changes during galvanostatic electrolysis. 3.2.1. Influence of Initial pH on Solution pH during Galvanostatic Electrolysis Using Aluminum Electrodes. Figure 11-a presents the changes of pH vs time during galvanostatic electrolyses of 0.1 M NaCl aqueous solutions of different initial pH values using aluminum plates as electrodes at a current density of 10 mA cm2. This figure shows that pH changes over time depend on the initial pH value of the electrolyzed solution. These results demonstrate that regardless of initial pH value, the pH evolution versus time follows two stages: a stage of rapid pH change over time and a stage in which pH remains almost constant. It is noteworthy that the curve of pH vs time is ascending for solutions of initial pH lower or equal to 9 and is descending in alkaline solutions (pH > 9). For acidic initial pH conditions, the increase in pH can be explained by the formation of OH ions from the reduction of H+ or H2O on the surface of the cathode.30,31 In these circumstances, it appears that hydroxide ions formed at the cathode were not completely involved in 13369


Industrial & Engineering Chemistry Research the formation of hydroxo-aluminum species. Rather, the excess of OH ions increased the pH of the medium. In contrast, for solutions with initial pH higher than 9, the decrease in pH could be attributed to the consumption of higher amounts of OH ions to form anionic hydroxo-aluminum species such as Al(OH)4 and Al(OH)52. It should be taken into account that in highly alkaline conditions, the amount of Al3+ ions obtained by chemical dissolution plays a significant role in the total amount of aluminum ions generated in situ. By the end of electrolysis, the majority of aqueous solutions stabilize at a pH value close to 9. As can be observed, the closer the initial pH is to 9, the more rapid the stabilization of pH. The stabilization of pH to a value close to 9 for all initial values of pH can be explained by a buffering effect of hydroxo-aluminum species that balances the quasi-static variation of the hydroxyl ions concentration through the formation of monomeric and polymeric complexes of aluminum hydroxides. Several acidbase pairs can be formed in the medium which buffer the pH to a value around pH 9. 3.2.2. Influence of Current Density on Solution pH during Galvanostatic Electrolysis Using Aluminum Electrodes. Figure 11-b presents the changes of pH vs time during galvanostatic electrolyses at different current densities of 0.1 M NaCl aqueous solutions using aluminum plates as electrodes at initial pH 9. No significant change of pH with time was observed using different current densities, and similar profiles were obtained for all applied current densities. At the start of electrolysis, the pH dropped by one unit, and then it increases with time until it reaches a value close to its initial value of pH 9) and eventually stabilizes. This insignificant change of pH caused by changing current density during galvanostatic electrolysis can be explained by the fact that increasing the current density in the studied range results only in an increase of both the rate of the aluminum anode dissolution and production of H2 at the cathode surface. 3.2.3. Effect of Supporting Electrolyte on Solution pH during Galvanostatic Electrolysis Using Aluminum Electrodes. Figure 11-c presents the variation of pH vs time during galvanostatic electrolyses of aqueous solutions containing different supporting electrolytes (NaCl, Na2SO4, or NaH2PO4) using aluminum electrodes at a current density of 6 mA cm2. Figure 11-c clearly indicates that the nature of the supporting electrolyte has a significant effect on the variation of pH with time during galvanostatic electrolysis. In the presence of NaH2PO4, pH increased continuously until it reached a value of 11.4. Whereas in the presence of NaCl or Na2SO4, pH decreased slightly at the beginning and then increased gradually until it was stabilized at 9.3 and 10, respectively. These variations are probably due to the nature and stability of aluminum complexes formed in each electrolyte. In particular, the presence of NaH2PO4 in the solution leads to the formation of stable complexes with Al3+ ions (Al(PO4), AlH(PO4)+, and AlH2(PO4)2+).3235 As such, a major part of Al3+ ions formed during electrolysis reacts with phosphate ions. This leads to an excess of OH available in the solution, which increases solution pH with time. It is notable that in the presence of Na2SO4 or NaH2PO4, pH of the solution stabilized at a value greater than pH 10, which facilitates the dissolution of Al(OH)3 into anionic hydroxoaluminum species such as Al(OH)4 and Al(OH)52 and the formation of these species makes flocculation and electroprecipitation difficult to occur. While in the case of NaCl, pH remains in the range between 8.3 and 9.3 throughout the experiment. In this pH range, amorphous aluminum hydroxide (Al(OH)3) is

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precipitated, which has a large surface area that is beneficial for coagulation by adsorption.3945

4. CONCLUSION The influence of supporting electrolyte, initial pH, and current density on anodic dissolution of pure aluminum during electrocoagulation process was evaluated. The main conclusions of this work can be summarized in the following points: - Potentiodynamic polarization experiments showed that aluminum corrosion is largely dependent on supporting electrolyte and initial pH conditions. The presence of NaCl or Na2SO4 shifted the corrosion potential toward more negative values, indicating that corrosion of aluminum is catalyzed by the presence NaCl or Na2SO4. However, in the presence of NaCl, pitting corrosion occurs at a much lower potential than that in the presence of Na2SO4. In contrast, the presence of NaH2PO4 increased the corrosion potential, which indicates that the presence of NaH2PO4 inhibits the corrosion of aluminum. EIS and potentiostatic anodic current transients confirmed that pitting corrosion kinetics is accelerated by pH increase and the presence NaCl. Morphological observations proved that large localized pitting corrosion was clearly observable on the surface of aluminum plate polarized in the presence of NaCl. However, in the case of Na2SO4, small uniformly distributed pitting corrosion was observed. In contrast, the aluminum plate polarized in the presence of NaH2PO4 was completely covered with a layer of AlPO4. - Galvanostatic electrolyses demonstrated that measured concentrations of aluminum were slightly higher than theoretical values calculated using Faraday’s law and the gap between them enlarged with the increase in specific electrical charge. Super-Faradic excess of dissolved aluminum is primarily related to the chemical dissolution of aluminum, which is more significant at highly alkaline conditions. These results indicate that the amount of dissolved aluminum is basically related to pH conditions. Changes of pH during galvanostatic electrolysis mainly depend on initial pH and the nature of supporting electrolyte. In the presence of NaH2PO4, the pH increased continuously until it reached a value of 11.4. Whereas in the presence of NaCl or Na2SO4, pH decreased slightly at the beginning and then increased gradually until it stabilized at 9.3 and 10, respectively. - For acidic initial pH conditions, pH increases during electrolyses; however, it decreases for solutions with initial pH higher than 9. These results can be due to the nature of hydroxo-aluminum species formed during anodic dissolution of aluminum. In acidic and neutral conditions, it appears that hydroxide ions formed at the cathode were not completely involved in the formation of cationic and neutral hydroxo-aluminum species and the excess of OH ions increased the pH of the medium. In contrast, the decrease of pH in highly alkaline conditions could be attributed to the consumption of higher amounts of OH ions to form anionic hydroxo-aluminum species such as Al(OH)4 and Al(OH)52. The stabilization of pH to a value close to 9 for slightly acid, neutral, and alkaline initial values of pH can be explained by a buffering effect of hydroxo-aluminum species that balances the quasi-static variation of the concentration of hydroxyl ions through the formation of monomeric and polymeric complexes of aluminum hydroxides. 13370



’ AUTHOR INFORMATION Corresponding Author

*Phone: +21675392600. Fax: +21675392421. E-mail: nasr. [email protected] and/or [email protected].

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Anodic Dissolution of Pure Aluminum during Electrocoagulation

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