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Chemical impurities of aluminum used in this work (ppm by mass). equipped with a return-flow cooler. SEM images were taken with the Environmental Scanning Microscopy instrument (XL 30, Philips, Netherlands). Two scenarios were investigated in which aluminum was in contact with static solutions (without external agitation) and dynamic solutions (agitated magnetically). The experiments were performed by contacting one disc of aluminum (0.75 g) with 200 mL of the desired solution. The volume of 200 mL was chosen to avoid the saturation of the solution and preventing a solubility-control phenomenon in the leachate. Al?* present in aqueous solutions was determined by reverse titration of zinc sulfate solution with EDTA using xylenol orange as indicator [14]. The data presented are an average of two tests replicates with an error of 5%. Aluminum dissolved was calculated in term of weight loss using the expression: Table 1

Table 1 Chemical impurities of aluminum used in this work (ppm by mass). equipped with a return-flow cooler. SEM images were taken with the Environmental Scanning Microscopy instrument (XL 30, Philips, Netherlands). Two scenarios were investigated in which aluminum was in contact with static solutions (without external agitation) and dynamic solutions (agitated magnetically). The experiments were performed by contacting one disc of aluminum (0.75 g) with 200 mL of the desired solution. The volume of 200 mL was chosen to avoid the saturation of the solution and preventing a solubility-control phenomenon in the leachate. Al?* present in aqueous solutions was determined by reverse titration of zinc sulfate solution with EDTA using xylenol orange as indicator [14]. The data presented are an average of two tests replicates with an error of 5%. Aluminum dissolved was calculated in term of weight loss using the expression: Table 1

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Fig. 1. Aluminum dissolved in HCI (a), NaOH (b), H2SO, (c) and HNO; (d) at 0.5, 1 and 2 M at 25 °C in static media and fitting the experimental data obtained with HCI (e) and NaOH (f) by heterogeneous kinetic models.
Fig. 1. Aluminum dissolved in HCI (a), NaOH (b), H2SO, (c) and HNO; (d) at 0.5, 1 and 2 M at 25 °C in static media and fitting the experimental data obtained with HCI (e) and NaOH (f) by heterogeneous kinetic models.
Fig. 2. SEM images of aluminum surface after contacting at 25 °C the solutions of HCI], NaOH, HNO3 and H2SO, (0.5 and 2 M) at different contact times.
Fig. 2. SEM images of aluminum surface after contacting at 25 °C the solutions of HCI], NaOH, HNO3 and H2SO, (0.5 and 2 M) at different contact times.
Fig. 3. Aluminum dissolved in 2 M HCI (a), 2M NaOH (b), 2 M H2SO, (c) and 2 M HNO; (d) at 40, 50, 60, 70 and 80 °C in agitated media and fitting the experimental dat obtained with 2 M HCI (e) and 2 M NaOH (f) by heterogeneous kinetic models. NaOH solution dissolved aluminum more rapidly at low tem- perature than HCI (Fig. 3b). In fact, at 40 and 50°C Al release at- tained 90.7 and 210.6 mg/cm? after 30 min of contact with NaOH while it reached 34.7 and 79.7 mg/cm? with HCl after the same period of time and temperatures. At 60 °C almost the same results were obtained in both cases and at 70 and 80°C, higher aluminum release was obtained with HCl compared to NaOH. Aluminum dis- solution was controlled by chemical reaction in both solutions. The rate constants calculated with HCl at 70 and 80°C (0.0141 and 0.0214 min~') were approximately 2 times higher than those ob- tained with NaOH (Fig. 3e and f) at the same temperatures (0.0069 and 0.0106 min~'). Thus, acidic media was more aggres- sive at high temperature than alkaline media and the contrary be- came true as the temperature decreased. Apparent activation energies were calculated using the experimental rate constants of the kinetic models fitted. The activation energy was determined to be 86.5 kJ/mol with HCl and 52.4 kJ/mol with NaOH. Reported Ea for the reaction between aluminum and NaOH was in range 42.5- 68.4 kJ/mol [23-25] and that for aluminum dissolved in HCl was found to be equal to 65.2-71 kJ/mol [12,26]. The high value of the activation energy obtained with HCl indicated a temperature dependent reaction which is consistent with the results obtained. In HNO3 and H2SO, the dissolution of aluminum remained very
Fig. 3. Aluminum dissolved in 2 M HCI (a), 2M NaOH (b), 2 M H2SO, (c) and 2 M HNO; (d) at 40, 50, 60, 70 and 80 °C in agitated media and fitting the experimental dat obtained with 2 M HCI (e) and 2 M NaOH (f) by heterogeneous kinetic models. NaOH solution dissolved aluminum more rapidly at low tem- perature than HCI (Fig. 3b). In fact, at 40 and 50°C Al release at- tained 90.7 and 210.6 mg/cm? after 30 min of contact with NaOH while it reached 34.7 and 79.7 mg/cm? with HCl after the same period of time and temperatures. At 60 °C almost the same results were obtained in both cases and at 70 and 80°C, higher aluminum release was obtained with HCl compared to NaOH. Aluminum dis- solution was controlled by chemical reaction in both solutions. The rate constants calculated with HCl at 70 and 80°C (0.0141 and 0.0214 min~') were approximately 2 times higher than those ob- tained with NaOH (Fig. 3e and f) at the same temperatures (0.0069 and 0.0106 min~'). Thus, acidic media was more aggres- sive at high temperature than alkaline media and the contrary be- came true as the temperature decreased. Apparent activation energies were calculated using the experimental rate constants of the kinetic models fitted. The activation energy was determined to be 86.5 kJ/mol with HCl and 52.4 kJ/mol with NaOH. Reported Ea for the reaction between aluminum and NaOH was in range 42.5- 68.4 kJ/mol [23-25] and that for aluminum dissolved in HCl was found to be equal to 65.2-71 kJ/mol [12,26]. The high value of the activation energy obtained with HCl indicated a temperature dependent reaction which is consistent with the results obtained. In HNO3 and H2SO, the dissolution of aluminum remained very
The results show that the mixture A1 dissolved less amount of aluminum than 2 M HCl as shown in (Fig. 4a). Thus, longer reaction time is needed to the mixture Al to reach the same results ob- tained with 2 M HCI. The mixture A2 dissolved the same amount of aluminum as 2 M HNO; as shown in (Fig. 4b). In fact, similar re- sults were obtained in both cases suggesting that the aggressive behavior of HCl on aluminum was suppressed by the presence of nitrates [27]. With the mixture A3 a blockage of the reaction was observed over the whole reaction time where the weight loss sta- bilized at 58.1 mg/cm? (Fig. 4c). With 2M H,SO, and HNO; at 80°C, 188.2 and 164.9mg/cm? of aluminum were dissolved respectively after 4.5 h. The dissolution was interrupted with their mixture probably because of the adsorption of sulfate or/and ni- trate ions leading to protect the solid surface. Thus, in all cases, an inhibiting effect was observed when acids were mixed, and this effect was more pronounced in the presence of HNOs. Increasing the concentration of the acids to 4M led to a strong promoting effect observed with the mixture B1 which dissolved rapidly aluminum within 5 min as shown in (Fig. 4a). 4M HCI dissolved totally aluminum after 30 min while 4M H2SO, dissolved only In the case of the mixture B3 the induction period was lowered to 90 min (4.5 h minimum with A3) after which the release of alu- minum increased gradually reaching 277.7 mg/cm? after 4.5h (Fig. 4c). Thus, the blockage registered with the mixture at 2M over the whole reaction time seems to be less persistent at 4 M. The kinetic study showed that aluminum release was controlled by chemical reaction with the mixtures A1 and B1 (Fig. 5a) while with the mixture HCl + HNO3 the mechanism controlling the disso- lution varied with the concentration of acids. In fact, with A2 trans- port through boundary layer controlled the dissolution over the whole reaction time while with B2 both mechanisms - chemical reaction in range 2-12 min and transport trough boundary layer in range 14-60 min - controlled the dissolution. Aluminum release was controlled by chemical reaction with the mixture B3 after the elapses of the induction period (induction periods were excluded from the kinetic calculations with all mixtures). It is worth noting that no induction period was observed with the mixtures B1 and B2 indicating that the barrier oxide was instantaneously dissolved. This was concluded from the results obtained since after only 1 min of contact with both mixtures high amounts of Al?* were present in the solutions. It should also be noted that increasing the concentration of both H2SO,4 and HNO; to 4 M did not increase Fig. 4. Effect of the mixture of acids: HCl + H2SO, (a), HCl + HNO3 (b) and H,SO, + HNO; (c) on aluminum dissolution at 80 °C.
The results show that the mixture A1 dissolved less amount of aluminum than 2 M HCl as shown in (Fig. 4a). Thus, longer reaction time is needed to the mixture Al to reach the same results ob- tained with 2 M HCI. The mixture A2 dissolved the same amount of aluminum as 2 M HNO; as shown in (Fig. 4b). In fact, similar re- sults were obtained in both cases suggesting that the aggressive behavior of HCl on aluminum was suppressed by the presence of nitrates [27]. With the mixture A3 a blockage of the reaction was observed over the whole reaction time where the weight loss sta- bilized at 58.1 mg/cm? (Fig. 4c). With 2M H,SO, and HNO; at 80°C, 188.2 and 164.9mg/cm? of aluminum were dissolved respectively after 4.5 h. The dissolution was interrupted with their mixture probably because of the adsorption of sulfate or/and ni- trate ions leading to protect the solid surface. Thus, in all cases, an inhibiting effect was observed when acids were mixed, and this effect was more pronounced in the presence of HNOs. Increasing the concentration of the acids to 4M led to a strong promoting effect observed with the mixture B1 which dissolved rapidly aluminum within 5 min as shown in (Fig. 4a). 4M HCI dissolved totally aluminum after 30 min while 4M H2SO, dissolved only In the case of the mixture B3 the induction period was lowered to 90 min (4.5 h minimum with A3) after which the release of alu- minum increased gradually reaching 277.7 mg/cm? after 4.5h (Fig. 4c). Thus, the blockage registered with the mixture at 2M over the whole reaction time seems to be less persistent at 4 M. The kinetic study showed that aluminum release was controlled by chemical reaction with the mixtures A1 and B1 (Fig. 5a) while with the mixture HCl + HNO3 the mechanism controlling the disso- lution varied with the concentration of acids. In fact, with A2 trans- port through boundary layer controlled the dissolution over the whole reaction time while with B2 both mechanisms - chemical reaction in range 2-12 min and transport trough boundary layer in range 14-60 min - controlled the dissolution. Aluminum release was controlled by chemical reaction with the mixture B3 after the elapses of the induction period (induction periods were excluded from the kinetic calculations with all mixtures). It is worth noting that no induction period was observed with the mixtures B1 and B2 indicating that the barrier oxide was instantaneously dissolved. This was concluded from the results obtained since after only 1 min of contact with both mixtures high amounts of Al?* were present in the solutions. It should also be noted that increasing the concentration of both H2SO,4 and HNO; to 4 M did not increase Fig. 4. Effect of the mixture of acids: HCl + H2SO, (a), HCl + HNO3 (b) and H,SO, + HNO; (c) on aluminum dissolution at 80 °C.
‘ig. 5. Fitting the experimental data obtained with the mixture of acids at 2 and 4M by heterogeneous kinetic models controlled by chemical reaction (A) and controlle nass transport through boundary layer (x).
‘ig. 5. Fitting the experimental data obtained with the mixture of acids at 2 and 4M by heterogeneous kinetic models controlled by chemical reaction (A) and controlle nass transport through boundary layer (x).
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Mineral acidsEffect of concentration
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