Study of ferrous iron oxidation in Morocco drinking water in an airlift reactor

Study of ferrous iron oxidation in Morocco drinking water in an airlift reactor

Available online at Chemical Engineering and Processing 47 (2008) 1877–1886 Study of ferrous iron oxidation in Morocco drinkin...

650KB Sizes 0 Downloads 11 Views

Available online at

Chemical Engineering and Processing 47 (2008) 1877–1886

Study of ferrous iron oxidation in Morocco drinking water in an airlift reactor N. El Azher a , B. Gourich a,∗ , C. Vial b , M. Belhaj Soulami c , M. Ziyad d a


Laboratoire de G´enie des Proc´ed´es, Ecole Sup´erieure de Technologie de Casablanca, BP 8012, Oasis Casablanca, Morocco Laboratoire de G´enie Chimique et Biochimique, Universit´e Blaise Pascal, 24 avenue des Landais, BP 206, F-63174 Aubi`ere Cedex, France c Laboratoire de G´ enie des Proc´ed´es et d’Environnement, ENIM BP 753 Agdal Rabat, Morocco d D´ epartement de Chimie, Facult´e des Sciences, Universit´e Mohammed V, avenue Ibn Battouta, BP 1014, Rabat, Morocco Received 23 April 2007; received in revised form 29 June 2007; accepted 14 October 2007 Available online 22 October 2007

Abstract Although ferrous iron removal from drinking water by aeration has been studied for a while, there is still an uncertainty on the performance and the physicochemical mechanisms of iron(II) oxidation [S.K. Sharma, B. Petrusevski, J.C. Schippers, J. Water. Supply Res. Technol. Aqua, 54 (2005) 239–247]. A possible reason is the autocatalytic effect of ferric hydroxide particles, but this assumption is never validated quantitatively in practice because this catalytic effect has been investigated only under batch laboratory-controlled conditions. In this work, iron(II) oxidation has been studied on synthetic waters in a 63 L split-rectangular airlift reactor, the hydrodynamics and the mass transfer properties of which were described previously [N. El Azher, B. Gourich, C. Vial, M. Belhaj Soulami, A. Bouzidi, M. Ziyad, Biochem. Eng. J., 23 (2005) 161–167]. Experiments were carried out both under semi-batch and continuous flow conditions. The kinetic parameters derived from the experiments were consistent with the literature both under unsteady and steady-state conditions. Experimental results showed that the airlift reactor allowed simultaneously good mixing, mass transfer and pH control despite the strong sensitivity of the oxidation kinetics to pH. Data confirmed that recycling about 50 mg/L of ferric hydroxide particles in a slurry phase could decrease drastically the time necessary to reach the minimum admissible concentration of iron(II). For example, residence time could be reduced by a factor six at pH 7.0 under steady-state conditions, which may avoid the need for a further pH increase. © 2007 Elsevier B.V. All rights reserved. Keywords: Aeration; Airlift reactor; Iron(II) removal; Iron oxidation kinetics; Water treatment

1. Introduction Groundwaters are often mildly acidic and devoid of dissolved oxygen. Consequently, soluble ferrous iron, denoted iron(II), is usually present in groundwaters by contact with rocks and minerals, either in dissolved mineral form, Fe2+ , or associated with various organic, mineral or chelating agents [1,2]. An additional source for dissolved iron in drinking water may be the pipelines through which water flows, as water with high salinity or acidity from dissolved carbon dioxide or other acids is corrosive to metal pipes [3]. As a result, the presence of iron is probably the most common water problem faced by consumers and water treatment professionals [4]. Although iron does not present a health hazard, contrary to heavy metals, it brings unpleasantness of an aesthetic and

Corresponding author. Tel.: +212 22 23 15 60/65; fax: +212 22 25 22 45. E-mail address: [email protected] (B. Gourich).

0255-2701/$ – see front matter © 2007 Elsevier B.V. All rights reserved. doi:10.1016/j.cep.2007.10.013

organoleptic nature. Indeed, iron gives a rust colour to water, which can stain plumbing and sanitary facilities or even food industry products [5]. For example, iron will cause reddishbrown staining of porcelain, dishes, utensils and even glassware. Similarly, clothing may become stained a brownish colour; soaps and detergents do not remove these stains and the use of chlorine bleach and alkaline builders can actually intensify the stains. Iron also gives a metallic taste to water, making it unpleasant for consumption; the taste of beverages, such as tea and coffee, may also be affected. Additionally, iron deposits can build up in pipelines, tanks or water heaters and softeners, which reduce water flow rates by increasing pressure drops in the water distribution network. This problem is also associated to increased energy costs, like pumping water through constricted pipes or heating water with heating rods coated with iron minerals, but also to equipment cost when water supply or softening equipment must be replaced. Iron can also be at the origin of corrosion in drains and sewers due to the development of microorganisms, the ferrobacteries [5]. Additionally, when these die and slough


N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

off, this usually causes unpleasant odors and tastes. More generally, biofilms in iron pipelines are known to increase the iron(II) content in water [6]. As a result, the guideline for maximum iron content in drinking water defined by the World Health Organization to avoid aesthetic and organoleptic problems is 0.3 mg/L for iron. This corresponds to the limit defined by Moroccan regulations in drinking water, while the maximum iron concentration defined by the European Union [7] on the quality of water intended for human consumption is 0.2 mg/L. In current practice, iron concentration is however usually decreased by water treatment to a maximum of 0.1 mg/L, especially to avoid the above-mentioned corrosion problems. A literature review indicates that the standard methods for iron removal used in water purification processes include mainly: • ion exchange on cationic exchange resins, bentonite or zeolites [8]; • the oxidation–precipitation of soluble iron(II) into insoluble iron hydroxides, especially the ferric hydroxide Fe(OH)3 , that can be removed either by filtration on sand as well as on dual media when iron(II) concentration is lower than 7 mg/L or by decantation at higher concentration [5]. The use of membrane processes based on reverse osmosis [9], nanofiltration [9,10] and ultrafiltration [11] has also been reported. The oxidation–precipitation way constitutes the most common method, but it admits several variants as a function of the oxidant: oxidation can be carried out using a solid phase (greensand, pyrolox, etc.), chemical additives in solution (hypochlorite, permanganate, etc.) or dissolved gas (oxygen, chlorine, ozone, etc.). For acid highly carbonated waters, neutralisation with Ca(OH)2 may also favor the precipitation of iron hydroxide and iron carbonate [4]. Aeration is usually recommended for oxidizing ferrous iron in waters exhibiting high iron concentration (higher than 5 mg/L), so that costs for chemicals can be avoided [12]. Sometimes, in situ groundwater treatment in which ferrous iron is precipitated in the underground is possible [13,14]. Ozonation or chlorination is required only when complexed iron forms difficult to precipitate are present, i.e. when a stronger oxidant than oxygen is required. Aeration can be improved through action of specific microorganisms [4,15–22]. A review on the specific advantages and limitations of biological iron removal has been published recently [23]. Indeed, even in biological processes, physicochemical and biological mechanisms take place simultaneously [15]. 2. Kinetic and engineering aspects of iron(II) oxidation The present paper is focused on the physicochemical oxidation of iron(II) in water. The stoichiometry of ferrous iron oxidation into ferric iron using oxygen may be summarized as follows 4Fe2+ + O2 + 4H3 O+ → 4Fe3+ + 6H2 O


4Fe3+ + 12OH− → 4Fe(OH)3


This set of equations shows that 1 mol of oxygen will oxidize 4 mol of iron(II) under ideal conditions (respectively 1 mg of oxygen will oxidize 7 mg of ferrous iron). Additionally, 2 mol of acidity are produced per mol of iron(II) converted into iron(III). This kinetics of ferrous iron oxidation was established by several authors who agreed roughly on the following expression [24–28]: r(Fe2+ ) = K[Fe2+ ][OH− ] [O2 ] 2


However, the kinetic constant K was reported to depend strongly on the ionic strength and on the dissolved counteranions, while the influence of alkalinity could be adequately explained by ionic strength effects [28,29]. In summary, K decreased when ionic strength increased and the rate constant was found to decrease as a function of anion in the order: ClO4 − , NO3 − , Cl− , Br− , I− , SO4 − . Similarly, autocatalytic effects were observed when the products of oxygenation were the unstable lepidocrocite ␥FeOOH [27,29], although the addition of ferric hydroxide up to 5 mg/L did not seem to alter significantly the reaction rate [24]. As a result, this autocatalytic effect is not clearly mentioned in recent works [4,30]. From a technological point of view, water can be saturated by oxygen using: • pressure-driven aeration towers in which compressed air is sparged into water in form of fine bubbles [30]. The use of static mixers for mass transfer enhancement has also been reported [4,31]. On the other hand, the use of a multiple orifice spray reactor has been described for highly concentrated solutions, such as 150 mg/L iron(II) aqueous solutions [32]. • gravity-driven aerators in which the circulation of raw water on open-air cascades or trays induces natural aeration without the need for external air supply. In this case, the oxidation–precipitation step can sometimes be directly carried out in filters, as illustrated in [15]. Between these two systems, iron removal may also be carried out in aerated filters, but with an extra-aeration [22,33]. This is particularly true for biological iron removal processes that are mainly based on adsorptive filtration, whereas physicochemical treatments are usually carried out in pressure-driven units followed by floc filtration [23]. The present work is focussed on waters exhibiting high levels of iron, such as groundwaters, typically between 5 and 20 mg/L, even though values up to 80 mg/L have been reported [34]. In this case, multiphase gas–liquid reactors similar to bubble columns constitute usually the best way to carry out aeration [30]. A particular class of bubble columns denoted airlift reactors has been used, as these provide several advantages, such as good mixing due to the enhanced liquid circulation, high mass transfer rates, compactness and low operating and maintenance costs due to the absence of internals [35]. The objectives of this work can be summarized as follows. Although the above-mentioned literature shows that ferrous iron removal from drinking water by aeration has been studied for a while, there is still an uncertainty on the

N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886


Fig. 1. Experimental set-up.

performance and the physicochemical mechanisms of iron(II) oxidation [15,23]. In particular, the performance of physicochemical and biological treatments has sometimes been reported to exhibit the same efficiency under the same conditions. This demonstrates that it is never clear which mechanism is dominant. These discrepancies may be due to the strong pH-dependency of iron(II) oxidation by air, but they are often attributed to the autocatalytic effect of ferric hydroxide particles. However, this assumption is never qualitatively validated and this effect has been investigated only under batch laboratory-controlled conditions. Actually, there is also a lack of kinetic data obtained in pilot-plant facilities, especially under steady-state conditions. In the present work, the influence of operation parameters, such as gas flow rate, pH and the initial concentration of iron(II), will be investigated first under batch conditions, but in a 63 L airlift reactor whose hydrodynamics and mass transfer properties have been already described in detail [36,37]. Then, the autocatalytic effect of ferric hydroxide particles will be studied under similar conditions. Finally, iron(II) oxidation under steady-state conditions will be examined and the kinetic data as well as the performance of the technological solution proposed in this work will be discussed and compared with the literature. 3. Materials and methods The oxidation of ferrous iron was studied in a splitrectangular airlift reactor working either under semi-batch or continuous flow conditions (Fig. 1). The reactor volume VL was approximately 63 L. It was made of plexiglas and consisted of a

square column (0.20 m × 0.20 m) of 2 m height, divided equally into a riser and a downcomer section by a plexiglas baffle. More details can be found in previous works [36,37]. Experiments were carried out at 20 ± 1 ◦ C to minimize the effect of temperature on oxidation kinetics [27]. The influence of operating conditions, such as pH (between 6.5 and 7.5), superficial gas velocity in the riser UG , initial concentrations of soluble ferrous iron [Fe2+ ]0 (between 5 and 20 mg/L) and insoluble ferric iron [Fe(III)]0 (between 5 and 50 mg/L) were investigated, as well as the residence time of the liquid phase in continuous flow conditions. pH, dissolved oxygen and ferrous iron concentrations were monitored over time. The pH of the medium was measured using a WTW pH-meter (model pH197i, WTW, Germany) with a resolution of 0.01 and maintained constant by the controlled addition of carbon dioxide CO2 . The dissolved-oxygen levels were recorded using a WTW oxygen-meter (model OXI197i). All sensors were connected to a data acquisition system. Compressed air and CO2 were injected into the reactor and controlled using calibrated rotameters (Emerson Brooks, model 1355). UG could be varied from 0.01 up to 0.085 m/s. Synthetic waters were prepared using deaerated tap water (the properties of which are reported in Table 1) in which the initial ferrous iron concentration [Fe2+ ]0 was adjusted by the addition of iron(II) sulphate (FeSO4 ·7H2 O) and pH was adjusted using H2 SO4 aqueous solutions. Iron(II) sulphate was chosen because sulphate anions are a common constituent of most groundwaters and acid mine waters [38]. Additionally, sulphate anions have already been reported to decrease the rate constant in Section 2; as a result, the synthetic water composition cannot be said to be more favorable to iron


N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

Table 1 Alkalinity (mg/L NaHCO3 ) Total hardness (mg/L CaCO3 ) Turbidity (NTU) Conductivity at 20 ◦ C (S/m) Chloride (mg/L Cl− )

100 350 0.15 0.16 392

oxidation by air than in most works from the literature. Nitrogen bubbling was used for water desoxygenation. The desired ferrous sulphate concentrations were added to the tap water only when dissolved oxygen levels were below 1.0 mg/L. For the study of the autocatalytic effect of ferric hydroxide, commercial Fe(OH)3 particles were dispersed in the synthetic iron(II)-rich waters to adjust the initial iron(III) concentration [Fe(III)]0 . In the continuous mode, the liquid flow rate was varied between 25 and 95 L/h using a peristaltic pump, which corresponds to residence time values of the liquid phase τ ranging between 40 and 150 min; in this case, [Fe2+ ]0 and [Fe(III)]0 correspond to the inlet concentration of soluble iron(II) and insoluble iron(III), respectively. For the measurements of iron(II) concentration in the experiments [Fe2+ ], samples were recovered from the centre of the downcomer in semi-batch conditions and from the outlet flow stream in continuous flow conditions. Then, they were mixed to a 25% H2 SO4 aqueous solution in order to limit further ferrous iron oxidation and stored for subsequent analysis. Measurements were carried out using a spectrophotometric determination with 1,10-phenanthroline which is effective even in the presence of large amounts of iron(III) [39]. 4. Results 4.1. Effect of air flow rate The evolution of the performance of iron removal as a function of the superficial gas velocity was studied in semi-batch conditions in order to define the optimum air flow rate. Experimental data showed that the signal from the dissolved oxygen sensor remained constant over time during the batch experiments and that its average value was rather close to that of the oxygen solubility C*, regardless of pH and gas flow rate. As a result, oxygen mass transfer is clearly not the limiting step of iron(II) oxidation in the range of pH studied (between 6.5 and 7.5): when oxygen concentration in the liquid phase reaches the saturation, it remains constant all along the experiment (data not shown). Similarly, the change in mixing conditions due to the increase of UG did not affect apparently the conversion yield. This is clearly confirmed by Fig. 2 which shows the evolution of the ferrous iron concentration with time of aeration for different superficial gas velocities. These results could be expected. According to many authors [30,40–42], the dimensionless Hatta number is defined as   r(Fe2+ ) DO2 Ha = (4) 4C∗ kL

Fig. 2. Effect of gas flow rate on iron(II) oxidation at pH 7.0.

using [Fe2+ ]0 = 20 mg/L, DO2 = 2.0 10−9 m2 /s, kL ≈ 4.10−4 m/s and Eq. (3) for the kinetics r(Fe2+ ) with [O2 ] = C* and [Fe2+ ] = [Fe2+ ]0 . Ha varies between 5 × 10−4 and 5 × 10−3 when pH is varied between 6.5 and 7.5. This corresponds to kinetically slow gas–liquid reactions in the film (Ha < 0.2), which means that reaction takes mainly place in the liquid bulk. However, to validate experimental data, the reaction must also be slow in the liquid phase. This can be checked for example by the estimation of the following ratio [42]: α =

(1 − εG )kL aDO2


in which εG is the gas hold-up. When Ha < 0.02, α < 1000 and that simultaneously residence/reaction time is high in comparison to the characteristic time of mass transfer, the actual oxygen concentration in the liquid phase remains close to the saturation concentration. In this work, α is about 100–200 and kL aτ > 20 on the basis of previous works [36,37]. As a conclusion, the reaction is slow both in the film and in the liquid phase. Similarly, Gourich et al. [37] have shown that mixing time in the split-rectangular airlift reactor is nearly independent of UG in the range studied in this work and that it is about 30 s. As reaction time is about 60 min in Fig. 2 and oxygen concentration in liquid phase is close to saturation concentration, the reactor can be considered as perfectly mixed, regardless of UG value. The spatial heterogeneity of bubble concentration between riser and downcomer that characterizes airlift reactors has therefore no significant influence on reaction rate. As a conclusion, a superficial gas velocity equal to 0.033 m/s has been chosen for all the following experiments: indeed, this corresponds to the beginning of the plateau region in the liquid recirculation velocity vs. UG curve, i.e. it maximizes liquid circulation at the lowest power requirements, while gas recirculation is minimized in the riser [37]. 4.2. Effect of pH The results of semi-batch experiments using [Fe2+ ]0 ≈ 10 mg/L, varying pH between 6.8 and 7.3 at UG = 0.033 m/s, are

N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886


Fig. 4. Evolution of the slope k as a function of [OH− ]2 when [Fe2+ ]0 ≈ 10 mg/L.

4.3. Effect of the initial iron(II) concentration

Fig. 3. Evolution of iron(II) concentration vs. time as a function of pH for [Fe2+ ]0 ≈ 10 mg/L: (a) evolution vs. time and (b) semi-logarithmic plot.

reported in Fig. 3a. These plots indicate that higher pH increases reaction rate, as expected from Eq. (3). Consequently, the time required to achieve 90% yield for pH 6.9, 7.1 and 7.3 are 43, 23 and 7.5 min respectively. Fig. 3b shows that the evolution of [Fe2+ ] with time follows an exponential decrease and can be assimilated to a first-order mechanism. This is in agreement with previous results from the literature. Using the results of Section 4.1 on oxygen concentration together with Eq. (3), the following expression can be derived for [Fe2+ ]: −

d[Fe2+ ] 2 = K[Fe2+ ][OH− ] [O2 ] ≈ k[Fe2+ ] dt

The experimental results obtained in semi-batch experiments for an initial concentration of ferrous iron equal to 5, 10, 15 and 20 mg/L respectively, keeping all other experimental conditions identical, are reported in Fig. 5. This shows that iron(II) oxidation is much faster when [Fe2+ ]0 increases. This behaviour is however incompatible with the first-order kinetics described by Eqs. (3) and (6). As a first-order mechanism was suggested in Section 4.2 when the influence of pH was studied at constant [Fe2+ ]0 value, faster oxidation rate with increasing iron(II) concentration may be attributed to the formation of substantial ferric iron in the medium as a result of oxidation through catalytic effect on the kinetic constant [28,30]. These authors have considered that a heterogeneous mechanism based on iron(II) adsorption on Fe(OH)3 particles coexists with the homogeneous mechanism described by Eq. (3). This corresponds to the following kinetics at constant pH when the assumption of kinetically


By fitting experimental data with Eq. (6), K and k values have been estimated from the slopes of Fig. 3b. Fig. 4 plots the evolution of k as a function of the concentration of hydroxide anion and shows that k is proportional to [OH− ]2 . This constitutes an experimental confirmation of Eq. (3). A quantitative relation between k and [OH− ] is also reported in Fig. 4 for modeling purpose to account for the influence of pH.

Fig. 5. Evolution of iron(II) concentration vs. time as a function of [Fe2+ ]0 at pH 7.0.


N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

three times higher due to autocatalysis. While α remains nearly constant and kL aτ > 20, the reaction remains slow both in the film and the liquid bulk despite the catalytic effect of ferric hydroxide particles [42], in accordance with the experimental data provided by the oxygen probe. 4.5. Determination of the kinetics of iron oxidation Experimental results of Section 4.3 can now be reanalyzed using the kinetic data of Section 4.4. The determination of k at pH 7.0 can be achieved through several methods. When no ferric hydroxide is present at t = 0 and [Fe2+ ]0 is low (≤5 mg/L), Eq. (6) should apply, which corresponds to [Fe2+ ] ≈ [Fe2+ ]0 exp(−kt) Fig. 6. Evolution of iron(II) concentration vs. time as a function of [Fe(OH)3 ]0 at pH 7.0.

slow gas–liquid reaction is assumed: d[Fe2+ ] − = (k + kS [Fe(III)])[Fe2+ ] dt



When no ferric hydroxide is present at t = 0, but [Fe2+ ]

0 is higher than 5 mg/L, Sung and Morgan [27] provided Eq. (10) that corresponds to the analytical solution of Eq. (7):

[Fe2+ ] =

[Fe2+ ]0 (k + kS [Fe2+ ]0 ) kS [Fe2+ ]0 + k exp[(k + kS [Fe2+ ]0 )t]


In this expression, kS is the kinetic constant of the heterogeneous reaction. The validity of Eq. (7) will be tested in the next section.

When both [Fe2+ ]0 and [Fe(III)]0 cannot be neglected, the analytical solution of Eq. (7) that was calculated in this work is

4.4. Effect of Fe(OH)3 addition

[Fe2+ ] = [Fe2+ ]0

Known concentrations of ferric hydroxide between 5 and 50 mg/L have been dispersed into a 5 mg/L iron(II) aqueous solution (pH 7.0). This iron(II) concentration was chosen to minimize autocatalytic effects due to ferric hydroxide formed by iron(II) conversion, so that [Fe(III)] could be assumed to be constant and nearly equal to the initial concentration [Fe(III)]0 . In this case, Eq. (7) may be roughly simplified as follows −

d[Fe2+ ] ≈ kS [Fe(III)]0 [Fe2+ ] = kcat [Fe2+ ] dt


The results of the semi-batch experiments are reported in Fig. 6. This figure shows that the rate of ferrous iron oxidation was significantly influenced by the presence of ferric hydroxide and increased with [Fe(III)]0 . With an initial concentration of 50 mg/L, the time needed to reach 0.3 mg/L could be reduced from 120 to 20 min. As a result, the presence of ferric hydroxide induces obviously a strong autocatalytic effect on ferrous iron oxidation. In order to validate Eqs. (7) and (8), the apparent kinetic constant kcat was obtained by fitting experimental data with Eq. (8). kcat values are plotted as a function of the initial concentration of ferric hydroxide to check whether it is proportional to [Fe(III)]0 . Fig. 7 confirms the proportionality between 5 and 50 mg/L of Fe(III) and gives kS = 7.0 × 10−3 ± 0.5 × 10−3 L/(mg min). If Hatta number (Eq. (4)) is used to analyse the data, Ha estimation using Eq. (8) for autocatalysis together with [Fe2+ ]0 = 10 mg/L and [Fe(III)]0 = 50 mg/L when pH 7.0 gives about 5 × 10−3 , while it should have been 1.5 × 10−3 without autocatalysis. Although ferrous iron oxidation remains a kinetically slow gas–liquid reaction, Ha number when t = 0 is about

k + kS ([Fe2+ ]0 + [Fe(III)]0 ) + [k + kS [Fe(III)]0 ]

kS [Fe2+ ]0

× exp[(k + kS ([Fe2+ ]0 + [Fe(III)]0 ))t] (11) Actually, the attempts to fit the whole set of experimental data at pH 7.0 with Eq. (11) by adjusting simultaneously k and kS failed, probably because of the strong nonlinearity of Eq. (11). An alternative method was used to validate the kinetic parameters when both the homogeneous and heterogeneous oxidation mechanisms coexist. This consists in injecting the results of Fig. 7, kS = 0.007 L/(mg min), in Eq. (11) before fitting experimental data with Eq. (11) by the optimization of k. Calculations provided the following estimation: k = 0.07 ± 0.02 min−1 . This result is in good agreement with Fig. 4 that gave k = 0.055 min−1 .

Fig. 7. Evolution of kcat as a function of [Fe(OH)3 ]0 .

N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

Fig. 8. Evolution vs. time of the reaction yield as a function of the residence time of the liquid phase under continuous flow conditions at pH 7.0 and [Fe2+ ]0 ≈ 10 mg/L.

A model able to describe ferrous iron oxidation at pH 7.0 is now established. 4.6. Continuous flow experiments For steady-state experiments, the column was continuously fed with fresh medium. The influence of the residence time of the liquid phase τ was studied between 40 and 150 min with the inlet ferrous iron concentration fixed either at 5 or at 10 mg/L, pH around 7.0 and superficial gas velocity in the riser between 0.016 and 0.033 m/s. Once again, superficial gas velocity did not influence the results (data not shown), as in Section 4.1, due to the slow kinetics of the reaction. Fig. 8 shows clearly that the autocatalytic effect described in batch experiments occurred also in steady-state conditions: indeed, the yield increased from 82% to 93% when the inlet concentration [Fe2+ ]0 was increased from 5 to 10 mg/L. As a result, the yield χ of ferrous iron oxidation under steady-state conditions can be expressed as a function of the residence time τ, as follows when the autocatalytic behavior can be neglected ([Fe2+ ]0 ≤ 5 mg/L) χ=

kτ 1 + kτ


while it is defined by the following expression that relates linearly χ to the ratio χ/[(1 − χ)τ] k + kS [Fe2+ ]0 χ =

χ (1 − χ)τ


when ferrous iron concentration [Fe2+ ]0 is higher than 10 mg/L with kS ≈ 7.0 × 10−3 L/(mg min). Experimental results are presented in Fig. 9. They show clearly the increase of oxidation yield when residence time is increased. Using the Levenberg–Marquart algorithm with Eqs. (12), (13) and kS = 0.007 L/(mg min) for the optimization of k, one finds k = 0.08 ± 0.01 min−1 , which is in agreement with the results of Section 4.5.


Fig. 9. Evolution vs. time of the reaction yield as a function of [Fe2+ ]0 under continuous flow conditions at τ = 85 min and pH 7.0.

4.7. Discussion A comparison of our results with the kinetic data obtained under laboratory controlled conditions shows good agreement. For example, Sung and Morgan [27] reported that k = 6.1 × 1013 [OH− ] (1/min) 2


while we found in Fig. 4: k = 5.5 × 1013 [OH− ] (1/min) 2


At pH 7.0, there is also a good agreement with the literature data if the differences in ionic strength, alkalinity and water composition (especially in terms of anions) are accounted for. Tamura et al. [29] reported k ≈ 0.05 1/min, while we found k between 0.05 and 0.08 1/min. Similarly, kS was shown to be respectively 0.005 L/(mg min) [27] and 0.0065 L/(mg min) [29] at pH 7.0, while kS = 0.007 L/(mg min) in this work. These results demonstrate that the assumption of a perfectly mixed reactor is totally justified and point out the good mixing conditions achieved despite the scale of the reactor (63 L). Indeed, the high sensitivity of the oxidation of ferrous iron to pH highlights the good ability to maintain a constant pH in the airlift reactor. This shows that the data obtained in this work can be used for scale-up purpose provided good mixing conditions are maintained. As a result, the split-rectangular airlift reactor that has been shown to exhibit good mixing and mass transfer properties in previous works appears to be a useful tool for carrying out iron removal in drinking water. On the basis of the above-mentioned results, substantial savings on reactor volume are possible. Indeed, autocatalysis can be enhanced by recirculating the ferric hydroxide particles formed by ferrous iron oxidation in a slurry phase. A rough estimation of the decrease in residence time from τ (without Fe(OH)3 recirculation) to τ 0 (with Fe(OH)3 recirculation) under steady-state


N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

be drastically decreased without the need to adjust the pH too far from neutrality, which would require a downstream neutralization stage. Although it is difficult to compare quantitatively several types of processes that usually work with groundwaters of different compositions, the advantages of the recirculation of ferric hydroxide particles over the common biological oxidation process can be summarized as follows:

Fig. 10. Evolution of the VL /VL0 ratio with [Fe(OH)3 ]0 under continuous flow conditions at pH 7, as well as constant [Fe2+ ]0 and τ values.

conditions can be obtained using the following relation: τ0 k ≈ τ k + kS [Fe(III)]0


which is plotted in Fig. 10 at pH 7.0 using experimental data on k and kS . This corresponds exactly to the ratio of the homogeneous oxidation rate over the total reaction rate taking both mechanisms into account. From a theoretical point of view, this ratio quantifies the decrease in reactor volume requirements that could be expected at constant yield and liquid flow rate, provided a fraction of the ferric hydroxide particles formed by iron(II) oxidation was recycled in the airlift reactor to maintain [Fe(III)]0 nearly constant. Similarly, it quantifies the possible increase of liquid flow rate at constant reactor volume, yield and [Fe(III)]0 . In other words, Fig. 10 shows the possible savings on equipment cost (reactor volume) allowed by the catalytic effect. From a practical point of view, process improvements by Fe(OH)3 recirculation can be achieved only if the hydrodynamic and mass transfer properties are not negatively affected by the change of residence time. For example, Fig. 10 shows that reactor volume for ferrous iron oxidation can be roughly reduced by a factor 6 when the ferric iron concentration is increased from 0 to 50 mg/L at pH 7.0: for χ = 93% (τ = 150 min in Fig. 9), one should get the same yield for τ 0 = 25 min when [Fe(III)]0 = 50 mg/L. In this case, Ha remains lower than 0.02 (see Section 4.6); α should not change significantly at constant superficial gas velocity because the contribution of liquid flow rate remains small in comparison to the overall liquid velocity; as a result, mass transfer conditions could be maintained provided kL aτ 0 > 20, which is achieved in this case because kL aτ 0 ≈ 30 [36,37]. As a conclusion, the kinetics of the physicochemical oxidation of ferrous iron by air can be adjusted not only by pH, but also by using heterogeneous catalysis based on a slurry phase containing Fe(OH)3 particles. Consequently, the conversion may be dramatically enhanced and the residence time required to reach the maximum admissible level of iron in drinking water can

• Even when biological oxidation is applied, abiotic ferrous iron removal is reported to represent up to 50% of total iron removal [15,22,23], which shows that the physicochemical mechanisms are never negligible. As reaction rate may be increased by about a factor 6 using the recirculation of ferric hydroxide particles, close performances may be achieved: the half-life of the reaction at pH 7 is more than 12 min for air oxidation without catalysis, 2 min with [Fe(III)]0 = 50 mg/L and about 1 min for biological oxidation [22]. • As illustrated by the example in Section 4.6, airlift reactors are versatile tools and are less sensitive to rapid and large variations in the flow rate and water quality than reported for biological treatments [23]. Similarly, efficient oxidation by air in the airlift reactor can be started very quickly, while a long maturation time is necessary before full efficiency is achieved with biological oxidation processes (50–60 days for a new reactor and about 5 days after a 2-month shutdown) [23]. However, the disadvantages of the airlift reactor are its residence time that will be always higher than biological filtration units [22] and the related equipment costs due to reactor volume. Further works should therefore be aimed at investigating Fe(OH)3 recirculation for situations in which biological treatments fail, such as groundwaters containing large amounts of NH4 + , H2 S and Zn [23]. 5. Conclusions Ferrous iron removal from drinking water has been studied using iron oxidation based on aeration in a 63 L split-rectangular airlift reactor. Experiments were carried out both under semibatch and continuous flow conditions on synthetic waters rich in iron(II). The effect of operating conditions was investigated, including the influence of gas flow rate, pH, as well as the initial concentrations of ferrous iron [Fe2+ ]0 and ferric hydroxide [Fe(III)]0 . The optimum gas flow rate was estimated on the basis of previous results on reactor hydrodynamics. Experimental data confirmed that the iron oxidation kinetics was strongly sensitive to the pH and that it exhibited an autocatalytic behaviour that played a key role in waters with high levels of ferrous iron, such as groundwaters, typically when [Fe2+ ]0 was higher than 5 mg/L. Experiments in semi-batch and continuous flow conditions gave results in agreement with the literature and the kinetic parameters found in semi-batch conditions could therefore be applied for iron removal in drinking water. As a result, the splitrectangular airlift reactor was shown to be adequate for iron(II) removal because it presents good mixing, mass transfer and pH control properties. Despite the catalytic effect, iron(II) oxida-

N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

tion remained a kinetically slow gas–liquid reaction both in the film and the liquid bulk when pH was around 7, but maintaining a constant [Fe(III)]0 value seemed to be able to avoid a further increase of pH. Finally, the controlled recirculation of ferric hydroxide particles formed by iron oxidation in a slurry phase seems to constitute a promising tool to promote savings on energy consumption and equipment costs in water treatment dedicated to iron removal, especially when biological treatments are not suitable or efficient [23]. Appendix A. Nomenclature

C* oxygen solubility in water (kg/m3 ) DO2 diffusion coefficient of oxygen (m2 /s) [Fe(III)] insoluble ferrous iron concentration (kg/m3 ) [Fe(III)]0 initial (semi-batch)/inlet (continuous) concentration of insoluble ferric hydroxide particles (kg/m3 ) 2+ [Fe ] soluble ferrous iron concentration (kg/m3 ) [Fe2+ ]0 initial (semi-batch)/inlet (continuous) soluble ferrous iron concentration (kg/m3 ) Ha Hatta number k apparent homogeneous kinetic constant (s−1 ) apparent heterogeneous kinetic constant (s−1 ) kcat mass transfer coefficient (m/s) kL kS heterogeneous kinetic constant (m3 /(kg min)) K kinetic constant (m9 /(kg mol2 s)) dissolved oxygen concentration (kg/m3 ) [O2 ] − [OH ] hydroxide anion concentration (mol/m3 ) r(Fe2+ ) rate of iron(II) oxidation reaction (kg/(m3 s)) t time (s) UG superficial gas velocity in the riser (m/s) Greek letters α dimensionless ratio of liquid volume to film volume εG gas hold-up χ reaction yield τ residence time (s) τ0 residence time with Fe(OH)3 recirculation (s) References [1] H. Roques, Fondements th´eoriques du traitement chimique des eaux, Tec&Doc-Lavoisier, Paris, France, 1990. [2] S. Pich´e, F. Larachi, Oxidation kinetics of iron(II) complexes of trans1,2-diaminocyclohexanetetraacetate (cdta) with dissolved oxygen: reaction mechanism, parameters of activation and kinetic salt effects, Chem. Eng. Sci. 61 (2006) 3452–3462. [3] P. Sarin, V.L. Snoeyink, J. Bebee, K.K. Jim, M.A. Beckett, W.M. Kriven, J.A. Clement, Iron release from corroded iron pipes in drinking water distribution systems: effect of dissolved oxygen, Water Res. 38 (2004) 1259–1269. [4] Degr´emont, Water treatment handbook/M´emento technique de l’eau, 10th ed., Lavoisier SAS, Paris, France, 2005. [5] WHO, Water Sanitation and Health (WSH): Guidelines for Drinking-Water Quality, 3rd ed., World Health Organization, Geneva, Switzerland, 2004. [6] M.J. Lehtola, T.K. Nissinen, I.T. Miettinen, P.J. Martikainen, T. Vartiainen, Removal of soft deposits from the distribution system improves the drinking water quality, Water Res. 38 (2004) 301–610.


[7] European Union Council Directive 98/83/EC of 3 November, On the quality of water intended for human consumption, Official Journal of European Communities 330 (1998) 32–54. [8] S.S. Tahir, N. Rauf, Removal of Fe(II) from the wastewater of a galvanized pipe manufacturing industry by adsorption onto bentonite clay, J. Environ. Manage. 73 (2004) 285–292. [9] R. Molinari, P. Arguiro, L. Romeo, Studies on interactions between membranes (RO and NF) and pollutants (SiO2 , NO3 − , Mn2+ and humic acid) in water, Desalination 138 (2001) 271–281. ´ [10] A. Lastra, D. G´omez, J. Romero, J.-L. Francisco, S. Luque, J.R. Alvarez, Removal of metal complexes by nanofiltration in a TCF pulp mill: technical and economic feasibility, J. Membr. Sci. 242 (2004) 97–105. [11] K.-H. Choo, H. Lee, S.-J. Choi, Iron and manganese removal and membrane fouling during UF in conjunction with prechlorination for drinking water, J. Membr. Sci. 267 (2005) 18–26. [12] J.M. Wong, Chlorination–filtration for iron and manganese removal, J. Am. Water Works Association 76 (1984) 76–79. [13] C.A.J. Appelo, B. Drijver, R. Hekkenberg, M. de Jonge, Modelling in situ iron removal from ground water, Ground Water 37 (1999) 811– 817. [14] S. Mettler, In situ removal of iron from ground water: Fe(II) oxygenation and precipitation products in a calcareous aquifer, PhD Thesis, dissertation, Swiss Federal Institute of Technology, ETH, Zurich, Switzerland, 2002. [15] G. Dimitrakos Michalakos, J. Martinez Nieva, D.V. Vayenas, G. Lyberatos, Removal of iron from potable water using a trickling filter, Water Res. 31 (1997) 991–996. [16] C.V. Tremblay, A. Beaubien, P. Charles, A.J. Nicell, Control of biological iron removal from drinking water using oxidation-reduction potential, Water Sci. Technol. 38 (1998) 121–212. [17] M. Nemati, S.T.L. Harrison, G.S. Hansford, C. Webb, Biological oxidation of ferrous sulphate by Thiobacillus ferrooxidans: a review on the kinetic aspects, Biochem. Eng. J. 1 (1998) 171–190. [18] A. Mazuelos, I. Palencia, R. Romero, G. Rodriguez, F. Carranza, Ferric iron production in packed bed bioreactors: influence of pH, temperature, particle size, bacterial support material and type of air distributor, Miner. Eng. 14 (2001) 507–514. [19] A. Mazuelos, R. Romero, I. Palencia, F. Carranza, F.J. Borjas, Oxygen transfer in ferric iron biological production in a packed-bed reactor, Hydrometallurgy 65 (2002) 15–22. [20] C.J.N. Dempers, A.W. Breed, G.S. Hansford, The kinetics of ferrous-iron oxidation by Acidithiobacillus ferrooxidans and Leptospirillum ferrooxidans: effect of cell maintenance, Biochem. Eng. J. 16 (2003) 337–346. [21] K.R. Blight, D.E. Ralph, Effect of ionic strength on iron oxidation with batch cultures of chemolithotrophic bacteria, Hydrometallurgy 73 (2004) 325–334. [22] I.A. Katsoyiannis, A.I. Zouboulis, Biological treatment of Mn(II) and Fe(II) containing groundwater: kinetic considerations and product characterization, Water Res. 38 (2004) 1922–1932. [23] S.K. Sharma, B. Petrusevski, J.C. Schippers, Biological iron removal from groundwater: a review, J. Water Supply: Res. Technol. Aqua 54 (2005) 239–247. [24] W. Stumm, G.F. Lee, Oxygenation of ferrous iron, Ind. Eng. Chem. 53 (1961) 143–146. [25] P.C. Singer, W. Stumm, Acidic mine drainage: the rate-determining step, Science 167 (1970) 1121–1123. [26] R. Jobin, M.M. Ghosh, Effect of buffer intensity and organic matter on the oxygenation of ferrous iron, J. Am. Water Works Association 64 (1972) 590–595. [27] W. Sung, J. Morgan, Kinetics and products of ferrous iron oxygenation in aqueous systems, Environ. Sci. Technol. 14 (1980) 561–568. [28] W. Davidson, G. Seed, The kinetics of the oxidation of ferrous iron in synthetic and natural waters, Geochim. Cosmochim. Acta 47 (1983) 67–69. [29] H. Tamura, K. Goto, M. Nagayama, The effect of ferric hydroxide on the oxygenation of ferrous ions in neutral solutions, Corros. Sci. 16 (1976) 197–207. [30] M. Roustan, Transferts gaz–liquide dans les proc´ed´es de traitement des eaux et des effluents gazeux, Tec&Doc-Lavoisier, Paris, France, 2003.


N. El Azher et al. / Chemical Engineering and Processing 47 (2008) 1877–1886

[31] D. Klein, R.D. Neufeld, Use of multiple orifice spray reactor to accelerate ferrous iron oxidation in acidic mine water, Mine Water Environ. 24 (2005) 186–193. [32] T.E. Ackman, R.L.P. Kleinmann, In line aeration and treatment of acid mine drainage, US Bureau of Mines, Report of Investigations 8868, USA, 1984. [33] B.-Y. Cho, Iron removal using an aerated granular filter, Process Biochem. 40 (2005) 3314–3320. [34] H. Andersson, J. Johansson, Iron removal from groundwater in Rakai District, Uganda: a minor field study, MSc Thesis, dissertation, Lule˚a University of Technology, Sweden, 2002. [35] Y. Chisti, M. Moo-Young, Airlift reactors: characteristics, applications and design consideration, Chem. Eng. Commun. 60 (1987) 195–242. [36] N. El Azher, B. Gourich, C. Vial, M. Belhaj Soulami, A. Bouzidi, M. Ziyad, Effect of alcohol addition in gas hold-up, liquid velocity and mass transfer in split-rectangular airlift bioreactor, Biochem. Eng. J. 23 (2005) 161–167.

[37] B. Gourich, N. El Azher, C. Vial, M. Belhaj Soulami, M. Ziyad, Study of hydrodynamics, mixing and gas–liquid mass transfer in a split-rectangular airlift reactor, Can. J. Chem. Eng. 84 (2006) 539–547. [38] S. Santos, R. Machado, M. Neiva Correira, J.R. Carvalho, Treatment of acid mining waters, Miner. Eng. 17 (2004) 225–232. [39] H. Tamura, K. Goto, T. Yotsuyanagi, M. Nagayama, Spectrophotometric determination of iron(II) with 1,10-phenantroline in the presence of large amounts of iron(III), Talanta 21 (1974) 314–318. [40] W.D. Deckwer, Bubble Column Reactors, J. Wiley & Sons, Chichester, UK, 1992. [41] J.M. Coulson, J.F. Richardson, Chemical Engineering, Chemical Reactor Design, Biochemical Reaction Engineering Including Computational Techniques and Control, vol. 3, Pergamon Press, Oxford, 1979. [42] A.A. Shaikh, U.A. Al-Mubaiyedh, On the bulk-liquid reaction in isothermal reactive gas absorption, Chem. Eng. J. 70 (1998) 65–70.