Chemical Engineering Journal 235 (2014) 37–45

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Chemical Engineering Journal

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Recovery of high-purity magnesium solutions from RO brines byadsorption of Mg(OH)2(s) on Fe3O4 micro-particles and magneticsolids separation

1385-8947/$ - see front matter � 2013 Elsevier B.V. All rights reserved.http://dx.doi.org/10.1016/j.cej.2013.09.014

⇑ Corresponding author. Tel.: +972 4 8292191; fax: +972 4 8228898.E-mail address: agori@tx.technion.ac.il (O. Lahav).

Orly Lehmann, Oded Nir, Michael Kuflik, Ori Lahav ⇑Faculty of Civil and Environmental Engineering, Technion, Haifa 32000, Israel

h i g h l i g h t s

� A modified process is introduced for separating Mg(II) from 1st stage SWRO brines.� The method is based on adsorption/dissolution of Mg(OH)2 on Fe3O4 micro-particles.� The process proved highly feasible from both engineering and economic standpoints.� Three high purity solutions (>97%) were produced: MgCl2, MgSO4 and Mg(HCO3)2.

a r t i c l e i n f o

Article history:Received 20 June 2013Received in revised form 1 September 2013Accepted 3 September 2013Available online 13 September 2013

Keywords:RO brineMagnetiteMg(II)Mg recoverySeawater

a b s t r a c t

A new approach is presented for cost effective recovery of Mg(II) from 1st stage seawater reverse-osmosisbrines (salinity: twice seawater concentration). The process is based on precipitation of Mg(OH)2(s) on thesurface area of self-synthesized magnetite (Fe3O4) micro-particles and magnet-assisted separation of thesolids-slurry from the Mg(II)-depleted brine. Once separated from solution, the solids slurry is subjectedto acidic conditions (pH�4–6) under which Mg(OH)2(s) is recovered as Mg(II) with the counter anionbeing either SO2�

4 , Cl� or HCO�3 , depending on the choice of strong acid used in the dissolution step.The magnetite solids are then used in the following adsorption cycle. This paper focuses on proof-of-con-cept of the suggested process and on defining ranges for the major process operational conditions (Fe3O4

particle concentration; pH range maintained during Mg(OH)2(s) dissolution step; determination of thefavorable solid-aqueous separation technique, etc.). Once defined, the chosen operational conditionswere applied and shown to result in three high purity (>97%) Mg(II) solution products at costs whichare comparable with equivalent commercial products.

� 2013 Elsevier B.V. All rights reserved.

1. Introduction

The global installed capacity of seawater reverse osmosis desa-lination (SWRO) is growing rapidly, amounting to �80 Mm3/d atthe end of 2012. Typically, SWRO processes operate at �50% recov-ery, i.e. �0.5 m3 of desalinated permeate and �0.5 m3 of concen-trated seawater brine are generated from 1 m3 of seawater feed.The retentate (brine) stream, containing approximately twice theseawater total dissolved solids (TDS) concentration is returned tothe sea. Yet, due to its high mineral content, the brine can andshould be considered a potential resource. Previous worksaddressed the utilization of desalination brines for recovery ofeconomically valuable constituents [1]. Particularly, the potentialof extracting Mg(II) from brines was addressed, owing to the highconcentration of this ion in SWRO brines (�2600–2800 mg Mg/L)

and its high market demand [2,3]. Elemental magnesium alloysare increasingly used as a lightweight metal for, e.g., the automo-bile/aviation and mobile computing (e.g. laptops, tablets, cellphones) industries. Likewise, dry magnesium salts and Mg(II) solu-tions are widely used in industrial, agricultural, medical and envi-ronmental contexts [4].

Several works have suggested cost-effective technologies forseawater-based magnesium separation, tailored for specific appli-cations. Birnhack and Lahav [5] applied ion exchange to enrichdesalinated seawater with Mg2+ ions. Drioli et al. [6] presented ananofiltration/crystallization process for the recovery of MgSO4

from desalination brines. Lahav et al. [7] used seawater concen-trates, obtained by nanofiltration, as a magnesium source for thecrystallization of struvite (MgNH4PO4) particles from a nutrientrich stream within wastewater treatment plants. All these newmethods, aiming at specific products or applications, have yet tobe commercialized. In contrast, production of magnesium alloysfrom seawater and natural brines is a well-established industrial

38 O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45

process. In this process MgCl2 salt is produced from the sourcesolution, purified, fused and subjected to electrolysis, for genera-tion of Mg(s) and Cl2(g) [8]. The most common methods used forthe initial separation of magnesium from seawater and brines aresolar evaporation and induced chemical precipitation. Solar evapo-ration is usually applied for processing highly concentrated brinessuch as Dead Sea and Great Salt Lake waters [9,10]. This method isbased on the fact that MgCl2 (�7.7 M) is more soluble than NaCl(�6 M), enabling the precipitation of the latter to obtain concen-trated MgCl2 or MgCl2–KCl solutions. Dry-warm climate and largearea evaporation ponds are a prerequisite for application of thismethod, and the composition of the separated Mg(II) solutiondepends almost exclusively on the composition of the source solu-tion. The chemical precipitation method is typically applied to lessconcentrated brines, such as seawater [8]. The initial Mg(II) separa-tion step for this process includes dosage of strong base to induceMg(OH)2(s) precipitation. CaO (produced on-site by calcination ofquarried calcite) is usually the cheapest strong base alternativefor this purpose. The precipitated Mg(OH)2(s) is normally separatedfrom solution by gravity settling/filtration. Large settling basinsand high surface-area filtration media are required due to theamorphous nature of the fresh precipitant formed, which is charac-terized by low density sub-micron particles to which water mole-cules are bound [2]. Mg(OH)2(s) separation is followed by rinsing(optional) and acid dissolution (typically HCl or H2SO4) to obtainthe required Mg(II) solution. Out of the widely practiced methodsdescribed, chemical precipitation is probably the most flexible.By deciding on the rinsing method and the type of acid used fordissolving the Mg(OH)2(g), the purity and composition of the sepa-rated Mg(II) solution is determined. However, application of thismethod for SWRO desalination brines is limited due to large foot-print required, often too expensive or unavailable in seashorelocations.

In this paper we present an improvement to the conventionalchemical precipitation process, which significantly reduces thehydraulic retention time needed for separating the precipitatedMg(OH)2(s) from the source solution, thereby minimizing footprintand land demand. The new method is based on adsorption ofamorphous Mg(OH)2(s) to the surface of magnetite (Fe3O4) micro-particles, followed by magnetic separation of the mixture fromthe bulk seawater brine. Fe3O4 has been widely used for adsorptionof both anionic and cationic species, as well metal hydroxides[11,12]. Dixon [13] measured the adsorption capacity of Mg(II)on magnetite as a function of pH and found an almost vertical in-crease at pH values higher than 10, at which Mg(OH)2 starts to pre-cipitate. Fresh precipitants of iron hydroxides are readily adsorbedon magnetite surfaces, as shown by Lahav et al. [14]. This tendencyhas been verified also in other metal hydroxides such as Co(OH)2

and Cu(OH)2 as shown by Petrick et al. [15] and Klas et al. [16].In the following sections the new process is described, proof of

concept laboratory-scale results are presented and a rough costanalysis, demonstrating the process to be promising for the gener-ation of MgCl2 and particularly MgSO4 and Mg(HCO3)2 solutionproducts, is shown.

1.1. Process description

The suggested process, shown in Fig. 1, starts by dosing strongacid (either HCl or H2SO4) to 1st stage RO brine to reduce pH downto pH4.3 (i.e. to the H2CO�3 equivalence point). Then, CO2 removal iscarried out in conventional stripping towers with the purpose ofremoving (practically) all the dissolved inorganic carbon from thebrine. Next, the de-carbonated brine flows into an adsorption reac-tor where it mixes with magnetite micro-crystals in a fashion thatmaximizes the contact area between aqueous and solid surfacesand promotes homogenous adsorption. A strong base (CaO or

NaOH) is simultaneously added to the mixed reactor, inducing pre-cipitation of Mg(OH)2(s), which adsorbs upon formation on themagnetite particles surface, while other solution species (ions,ion pairs) remain predominantly in the dissolved form (note thatCaCO3 cannot precipitate since the total carbonate species concen-tration, CT, tends towards zero due to the previous CO2 strippingstep). The Ksp of Mg(OH)2(s) controls the adsorption pH, which re-mains constant at �pH10. In the next step the solids’ slurry(Fe3O4 particles coated by Mg(OH)2(s)) is separated from the brine,first by applying a magnetic field and gravity decantation of the Mgdepleted brine solution back to the sea, and then by applying vac-uum on the remaining solids slurry, with the aim of minimizing thepresence of background solution species (originating from residualbrine left in the slurry after the gravity decantation). The separatedmagnetite-Mg(OH)2(s) mixture may be further rinsed (optional),depending on the purity required from the ultimate Mg(II) solutionproduct. The desorption of Mg(II) from the magnetite surface areais carried out by addition of an acidified solution to the solid slurry.The nature of Mg(II) solution formed in this step is a function of thetype of acid used while the concentration is determined by the ap-plied acid volume. For example, MgCl2, MgSO4 and Mg(HCO3)2

solutions may be produced by using HCl, H2SO4 or CO2, respec-tively. Following the desorption step, the regenerated magnetiteis again magnetically separated from the Mg(II) solution, and recy-cled to the adsorption step. The CaO mass required to serve ascheap strong base in the process is produced onsite by CaCO3

calcination, using a standard lime kiln, as performed in the conven-tional industrial process [2]. The concentrated CO2 gas stream, alsoproduced in the calcination process, may be recycled to the desorp-tion step if Mg(HCO3)2 solution is the desired product, as per-formed in the soda-ash process described in [2]. Assuming thatonly a small amount of CO2 is lost, the recycled CO2 constitutesalmost 50% of the acid required for the Mg(OH)2 dissolution.

2. Materials and methods

2.1. Experimental system

All the experiments described in the paper were conducted in a1 L glass beaker with 1st stage SWRO brine obtained from theAshkelon desalination plant in Israel ([Mg2+]�2600 mg/L; alkalin-ity = �260 mg/L as CaCO3). The magnetite beads used in the studywere synthesized in our laboratory at 90 �C according to the proto-col shown in [16]. The particle size distribution of the synthesizedFe3O4 particles was measured (Fig. 2), and the most prevalent sizewas found to be 8.17 lm. This average size was considered to bethe ‘‘operational particle size’’, representing a cluster of particlesrather than a single one.

A Neodymium 8 � 8 � 1.5 cm magnet was attached to the bot-tom of the test cell to produce an external magnetic field, when re-quired for enhancement of solids separation.

2.2. Theoretical calculations

Theoretical calculations related to strong base dosages requiredin the Mg(OH)2(s) precipitation step, as well as determination ofstrong acid dosages required in the Mg(OH)2(s) dissolution stepwere performed using the PHREEQC software [17]. The Pitzerapproach for concentrated solutions was applied by the use of‘‘pitzer.dat’’ database, embedded in the program.

2.3. Experimental procedures

All the experiments reported in this paper consisted of the samemain steps: (1) as a preliminary step, the SWRO brine was acidified

Fig. 1. Schematic sequence of the proposed Mg(II) separation process from SWRO brines.

Fig. 2. Particle size distribution of exposed magnetite particles @pH9.5.

O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45 39

and CO2 was stripped from it to prevent CaCO3 precipitation in thefollowing, basic pH, step. Over 99% of the ionic carbonate speciesconcentration was eliminated from the brine by dosing H2SO4 toattain pH�4.3 (at which the carbonate system comprises CO2(aq)

as, practically, the sole carbonate species); this was followed by30 min of intense air bubbling for CO2 stripping; (2) Mg(OH)2(s)

precipitation onto the magnetite particles surface: magnetite mi-cro-particles were added to the brine, which was then slowlydosed with a strong base solution while the brine was rapidlymixed at 200 RPM, resulting in gradual elevation of the pH valueand allowing the formed Mg(OH)2(s) to precipitate almost evenlyonto the surface of the Fe3O4 and Mg(OH)2-coated-Fe3O4 particles.In all the experiments described in the paper, for convenience rea-sons, KOH was used as the strong base although the most costeffective base is CaO, as mentioned in the Introduction section;(3) a solids-aqueous phase separation step: Once Mg(OH)2(s) hadprecipitated (30 min of mixing at 200 RPM were applied), the mag-nesium depleted brine was separated from the Fe3O4/Mg(OH)2

slurry. Most of the brine was drained gravitationally out whileapplying a magnetic force from the bottom, for retaining the solidsat the bottom of the beaker. In order to reach a ‘‘purer’’ solutionproduct (i.e. less foreign ions, manifested by a higher Mg/Na molarratio in the product solution), further separation techniques wereapplied to the Fe3O4/Mg(OH)2 slurry (retained in the beaker) suchas vacuum draining (via GF/A filters, sieve 1.6 lm), and a rinsingstep with RO permeate water. These steps further separated the

Fe3O4/Mg(OH)2 particles from the dissolved species left over fromthe magnesium-depleted brine; (4) Re-dissolution of the Mg(OH)2

precipitant to obtain the final product solution: The Fe3O4/Mg(OH)2 particles left in the beaker (in practical terms the wholeoriginal solids mass) were then immersed in RO permeate waterand dosed with acid (H2SO4, HCl or CO2 according to goal Mg-solu-tion products) until pH dropped to between pH6 and pH4, at whichconditions all the Mg(OH)2(s) rapidly re-dissolved into solution. Allthe required acid was dosed once (in order to attain a highMg(OH)2(s) dissolution rate) and then the solution was rapidlymixed for 15 min in order to verify that pH remained constant,indicating that all the Mg(OH)2(s) had indeed re-dissolved intosolution. Throughout the dissolution step the pH was not allowedto drop below pH�4 in order to avoid dissolution of the magnetiteparticles, which was reported to occur at a measureable rate onlyat much more acidic conditions (pH < �2) [14].When CO2(g) wasused as the acidifying agent, it was bubbled in excess into the testcell for 30 min. In this case, pH did not go below pH�6. Note thatthe range pH 4–6 was chosen for convenience reasons, asMg(OH)2(s) can be dissolved (although at a slower rate) at a pH va-lue as high as �9.

In each set of experiments one parameter was modified, withthe aim of testing its effect on the process and the characteristicsof the final product solution attained. The experiments were aimedat: (1) Determining process feasibility; (2) Defining an appropriaterange of operational conditions; and (3) Applying the process un-der the most appropriate operational conditions to produce threedifferent magnesium product solutions by using three differentacids. Table 1 lists the experimental conditions applied in each ofthe experiments.

In order to determine the settling rate of Mg(OH)2(s) in the pres-ence of various Fe3O4 concentrations, with and without applicationof a magnetic force, the SWRO brine was mixed with magnetitebeads resulting in different Fe3O4 particle concentrations of 0 to30 gFe/L brine and dosed with KOH (operational parameterspresented in Table 1). Following the dosage of the strong basethe solids slurry was allowed to settle for 30 min either in the pres-ence of an external magnetic force, or in its absence (i.e. gravitysettling). The volume occupied by the solids slurry was recordedafter 30 min and (Sludge Volume Index) SVI values were calculatedaccording to standard [18] and presented in units of ml/g solids,where ‘‘g solids’’ indicated the combined dry mass of Mg(OH)2(s)

and Fe3O4(s).

Table 1Summary of experimental conditions.

Fe3O4

conc.KOHdosage

Mg(OH)2(s)precipitated

Separation method of Fe3O4/Mg(OH)2 from Mg(II) depleted brine Acid agent

Magnetgravitydraining

Vacuumdraininga

Rinse step Vacuumdraining

Magnetgravitydraining

Units

Experiment goal gFe3O4�Fe/L

Molar gMgðOHÞ2�Mg/L

ml/g RO waterto dry Fe3O4

Determining process feasibilityTesting Mg(II) recovery efficiency 30 0.214 2.55 + + � � � Re-dissolved into

100 ml DI waterwith H2SO4 (pH�4)

Determining Mg(OH)2(s) settling rate inthe presence of various Fe3O4 conc.,with/without magnetic force

0, 10, 20and 30

0.214 2.55 Magnetgravity/gravityonly

� � � � NA

Determining appropriate range of operational conditionsEffect of Fe3O4 conc. on Mg(OH)2(s)

separation efficiency and productsolution purity

10, 20and 30

0.214 2.55 + + � � � Re-dissolved into100 ml DI waterwith H2SO4 (pH�4)

Effect of rinsing water volume onproduct solution purity

10 0.198 2.36 + + 10, 20,30 and40 (15 min,mixed rapidly)

+ � 100 ml RO waterwith H2SO4 (pH�4)

Application of the processProducing three different Mg(II)

product solutionsb10 0.198 2.36 + + 20 (15 min,

mixed rapidly)+ � 50 ml RO water;

H2SO4/HCl/CO2(g)

a 2 ml Fe3O4/Mg(OH)2 slurry samples were vacuum drained using GF/A filters (sieve 1.6 lm), decreasing the mass water content to 50% (water content was measured bydrying the samples at 900C for 24 h).

b These experiments were carried out using the experimental conditions found to be favorable in previous experiments.

40 O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45

2.4. Analyses

Coupled plasma-atomic emission spectrometry (ICP-AES) wasused to determine Mg(II), Ca(II), K(I), Na(I), B, Fe(II) and SO4(-II)concentrations. Cl(-I) concentration was measured using theargenometric method, according to Standard Methods [18].

Note that since ion-pairing effects are significant in seawaterbrines, the terminology Mg(II) (for example) was used throughoutthe manuscript to denote the overall magnesium concentration, i.e.free Mg2+ plus ion pairs such as MgSO0

4, MgCl+, etc.

3. Results and discussion

This section is divided in four. The first part was aimed at prov-ing feasibility and demonstrating the advantages of the processover currently employed Mg(II) separation processes. The secondpart addresses the required range of operational conditions; inthe third part the process is applied to result in three differentMg(II) product solutions, i.e. three different acidic substances wereused in the dissolution step and both the Mg(II) recovery efficiencyand the purity of the obtained product solutions are reported.Finally, a cost assessment is presented for the three productsolutions and compared with the cost of alternative commercially-available chemicals.

3.1. Feasibility tests

3.1.1. Determining the required strong base dosage for Mg(OH)2

precipitationTo determine the actual KOH demand per Mg mass precipitated

(in eq/eq units) a titration curve was recorded (Fig. 3; [KOH] dosedat 2.66 N) for the precipitation of Mg(OH)2(s) in the actual SWRObrine used in the experiments. To verify the empirical results, atheoretical calculation was also performed, using the aquaticchemistry software PHREEQC [17], to establish the required KOH

masses required for precipitating different Mg(OH)2 concentra-tions from the brine. Fig. 3 shows that at a dosage of �218 meqKOH/L brine (�pH11.0, arrow indicated point), a sharp increasewas observed in pH, indicating that all the Mg(II) present in thebrine had precipitated. The results of this titration curve corre-sponded very well with a theoretical PHREEQC simulation con-ducted with similar data.

3.1.2. Mg(II) mass balance and recovery efficiencyA mass balance was performed on Mg(II) in order to show that

practically all the Mg(II) that precipitated as Mg(OH)2 settled in-deed on the magnetite surface and was thereafter recovered inthe product water with very little losses incurring as a result ofthe applied rinsing and separation processes as described inSection 3.2.2. The results of the mass balance showed that over99% of the Mg(II) that precipitated was recovered into the productsolution conforming that hardly any Mg(II) mass was lost through-out the different steps of the process.

The efficiency of Mg(II) recovery within the process was calcu-lated as a percentage of the actual Mg(II) mass recovered in theproduct solution out of the Mg(II) recovery goal, determined bythe mass of strong base (KOH) initially dosed to the carbonate-depleted raw RO brine. To determine the Mg(II) recovery efficiencyvalues, experiments aiming at recovering 2.55 gMg/L brine, (KOHdosage of 0.214 M) were first conducted with magnetite particlesconcentrations of 30 gFe3O4-Fe/L (experimental conditions arepresented in Table 1). Since almost all of the precipitated Mg(II)was recovered, Mg(II) recovery efficiency could be also estimatedby comparing the initial Mg(II) concentration in the brine to theMg(II) concentration in the Mg(II) depleted brine. In these experi-ments a high Mg(II) recovery (�90%) was recorded.

According to theoretical calculation performed with PHREEQCthis KOH dosage (0.214 M) precipitates 2.55 gMg/L, i.e. �98% ofthe Mg(II) present in the raw RO brine. In practice only �90% ofthe Mg(II) was recovered in the process. The 8% difference may

O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45 41

be ascribed to either small inaccuracies in ion pairs equilibriumconstants used within the theoretical model or human errors(analytical, process-related – such as loss of Mg colloids duringthe solids-aqueous separation procedure; technical errors) inconducting the Mg(II) mass balances (or a combination of either).

3.1.3. Determining the settling rate of Mg(OH)2(s) in the presence ofvarious Fe3O4 particle concentrations

The major difference between the presented method and previ-ous methods for Mg(II) recovery from seawater/brines [3,19] is theFe3O4 addition. Fe3O4 has excellent settling characteristics (e.g.Lahav et al. [14]) attributed to its high specific density. SVI valuesin the range of 3–8 ml/g were reported for gravitational settling ofpure magnetite without the presence of a magnetic force (Klaset al. [20], Morgan et al. [21]). The settling efficiency can be en-hanced at the presence of a magnetic force. Adsorption of metalhydroxide solids on magnetite surfaces was shown by [15], [20]and [21] for Fe(OH)2, Ni(OH)2 and Zn(OH)2. To quantify the settlingrates of Mg(OH)2-coated Fe3O4 particles, solid-aqueous separationexperiments were conducted with various Fe3O4 particle concen-trations in the range 0 to 30 gFe/L (experimental conditions listedin Table 1). In these experiments 99% of the soluble Mg(II) presentin the RO brine was calculated to be precipitated as Mg(OH)2 by astoichiometric KOH dosage. In reality, Mg(II) recovery efficiency of�90% was attained out of the theoretical goal.

Following the precipitation, the combined solids slurry(Mg(OH)2 and Fe3O4) was allowed to settle for 30 min (both withand without external magnetic force), after which its volume wasrecorded. The resulting SVI results (in units of ml sludge per g sol-ids, after 30 min of settling) are shown in Table 2. Table 2 showsthat the magnetite particles covered with Mg(OH)2(s) settled veryrapidly under gravity conditions (no magnetic force applied), occu-pying a low volume, and allowing for easy draining of >90% of theMg-deficient RO brine. The high specific weight of the Fe3O4

particles (even when coated with the much lighter Mg(OH)2

precipitates) was responsible for the rapid settling. Moreover, inthe presence of a magnetic field the slurry settled even morerapidly as manifested by the lower SVI values recorded.Conversely, in the absence of Fe3O4 particles, Mg(OH)2(s) formedcolloidal particles which hardly settled at all. In the conventionalseparation method the low Mg(OH)2 settling rates in seawaterare considered the bottleneck of the technology [22] and give riseto design difficulties in industrial applications [2]. To overcome thepoor settling performance some authors suggested adding excessCa(OH)2 as strong base in the precipitation reaction, as well asrecirculating part of the treated sludge thereby adding nucleationsites [23].

Fig. 3. Titration curve for the precipitation of Mg(OH)2(s) from Ashkelon’s SWRObrine. (KOH concentration = 2.66 N).

The lowest SVI value shown in Table 2 was recorded with thehighest Fe3O4 concentration (30 gFe/L) both in the presence andabsence of magnetic force. Nevertheless, in all the experiments inwhich Fe3O4 particles were present the solid slurry settled excep-tionally well, and even better so when a magnetic force was ap-plied. From Table 2 it can be generally concluded that the use ofFe3O4 to absorb Mg(OH)2 solids upon their precipitation allowsfor efficient separation of Mg from RO brines both in the presenceand absence of external magnetic force.

More specifically, the settling characteristics of the slurry, asrepresented by the SVI values, were clearly affected by theFe3O4(s):Mg(OH)2(s) mass ratio. A low ratio (i.e. lower Fe3O4 massversus a constant Mg(OH)2(s) precipitated mass) led to thickerMg(OH)2(s) coating on the magnetite surface, both decreasing thespecific weight of the combined solids slurry and also reducingthe attraction of the coated Fe3O4 particles to the magnet, leadingto higher SVI values in all precipitation experiments. Since in Table 2the lower SVI values corresponded with high Fe3O4 concentrations,applying a high Fe3O4 concentration appeared to be advantageous.However, high solids concentrations in the reactor also result in lar-ger solids slurry volumes (per a given mass of separated Mg(II)),which correspond with a higher background brine water volumetrapped in it, leading eventually (following the dissolution of theMg(OH)2(s) in the final step) to less homogeneous (i.e. lower purity)Mg(II) solutions. This conflict was further addressed inSection 3.2.1.

3.1.4. Reactivity and physical endurance of the magnetite particlesover time

The efficiency of the process during multiple operation cycles(n = 10, experiments lasting over three weeks) was assessed byboth monitoring Mg(II) separation efficiency and measuring dis-solved Fe concentrations in the product solution, as an indicationof the resistance of the Fe3O4 particles over the cycles and possiblerelease of Fe(II) or Fe(III) to the product solution. Since the rawSWRO brine was characterized by very low dissolved iron(<0.2 mg/l) concentration, any iron measured in the product solu-tion could have resulted only from erosion of the magnetite parti-cles. The results showed that at the pH range at which the Mg(OH)2

dissolution was carried out (pH�4 to pH�6), Fe3O4 dissolution/erosion was insignificant. A very low concentration of up to0.05 mg Fe/L was measured in the product solution. This is a verylow value in comparison to the 30,000 mg Fe/L that were presentin solution as solid particles during the Mg(OH)2(s) dissolution step,resulting in [Mg]/[Fe] molar ratios higher than 100,000 in the prod-uct solution. With regard to Mg(II) recovery efficiency over multi-ple operation cycles, the results did not indicate a decline inrecovery efficiency with the progressing cycles (10 cycles weredone with the same Fe3O4 particles). These two aspects, however,should be further assessed over a prolonged period of time and lar-ger number of cycles in order to determine the required annualmagnetite renewal percentage.

3.2. Determining the most appropriate range for main processoperational conditions

3.2.1. Effect of the magnetite particles concentration on the Mg(OH)2(s)

separation efficiency and product solution homogeneityAs mentioned in Section 3.1.4, the decision on the most ade-

quate Fe3O4 concentration range poses a conflict. On the one hand,high Fe3O4 concentrations result in a higher surface area availablefor Mg(OH)2(s) adsorption/precipitation and thus on a thinner (aver-age wise) Mg(OH)2 layer on the magnetic particles, resulting in onlysmall decrease in both the specific weight of the particles and theirmagnetic attraction, and thereby in efficient settling. The largesurface area also minimizes the phenomenon of homogeneous

Table 2SVI of Fe3O4/Mg(OH)2 SWRO brine slurries with and without external magnetic force, at various initial magnetite concentrations.

Fe3O4 concentration (gMagnetite-Fe/Lbrine) 10 20 30

SVIa

Without magnet (ml/gsolids) 23.2 ± 0.72 13.5 ± 0.47 9.09 ± 0.20With magnet (ml/g solids) 6.68 ± 0.96 3.82 ± 0.31 3.62 ± 0.39

a ‘‘g solids’’ Indicate the combined mass of Mg(OH)2(s) and Fe3O4(s).

42 O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45

precipitation (i.e. precipitation in solution and not on the particles’surface area) and, as an outcome, increases the Mg(II) recovery effi-ciency. On the other hand, when the applied Fe3O4 concentration ishigh, the absolute volume that the settled solids-slurry occupies ishigher, resulting in higher background RO brine trapped in it, whichreduces the Mg(II) product water purity due to inclusion of a highermass of foreign ions in it.

In order to determine the most advantageous Fe3O4 concentra-tion range, experiments were conducted with various concentra-tions, with an ultimate aim of determining the highest productsolution purity. Fig. 4 shows the Mg/Na molar ratio measured inthe product solution in experiments conducted with magnetiteconcentrations of 10, 20 and 30 gFe/L, operated with the aim to re-cover 2.55 gMg/L brine and with H2SO4 as the acid used in the dis-solution step. The molar Mg(II) to Na(I) ratio in the productsolution (denoted Mg/Na) was chosen to serve as a measure ofthe purity of the Mg-rich product solution since Na(I) is presentat a very high concentration in the brine (>20 g/L) and is not addedexternally in the process. In order to maximize the product solu-tion purity, vacuum filtration was applied to the Fe3O4/Mg(OH)2(s)

solids slurry following the magnet-assisted gravity draining of theMg(II)-depleted brine. An average Mg(II) recovery efficiency of89.6% ± 0.62 of the recovery goal was achieved in all experiments.As expected, the results (Fig. 4) showed a slightly higher Mg/Na ra-tio in the product solution in the experiments conducted with thelower magnetite concentration (10 gFe/L), in which the volume ofthe Mg(II) depleted brine that was not drained was smaller, result-ing in higher product solution purity. This trend, i.e. that productpurity improves at lower magnetite concentrations is in contrastto the magnetic separation ability that is negatively affected bythe decrease in magnetite surface area available for Mg(OH)2(s)

adsorption (see Section 3.1.4). Further results (not shown) indi-cated that applying a magnetite concentration lower than 10 g/Lmay somewhat further improve product purity but would signifi-cantly detract from the magnetic separation ability of the particlesunder the applied magnetic field. To conclude, since Mg(II) separa-tion was reasonable at 10 gFe3O4-Fe and solution purity improved,this particle concentration was recommended by us for applicationin the process.

3.2.2. Effect of rinsing water volume on the purity of the productsolution

In preliminary experiments (results not shown) a significantimprovement was found in the product Mg/Na molar ratio (indica-tion of product solution purity) upon rinsing of the Fe3O4/Mg(OH)2

slurry with RO water, following the initial magnetic solid-aqueousseparation step. Fig. 5 shows the product Mg/Na molar ratio at-tained in experiments aimed at determining the preferable ratiobetween rinsing water volume and vacuum drained Fe3O4 masscoated with Mg(OH)2. The rinsed solids slurry was vacuum-drained again before being transferred to the final acid-dissolutionstep. Fig. 4 shows that the purity of the product solution improvedsignificantly by applying the rinsing step; however the efficiency ofapplying rinsing water almost leveled off beyond the addition of30 ml of rinsing water per 1 g of vacuum drained solids slurry. Thisresult was attributed to a certain very small adsorption of Na(I) on

the magnetite particles surface, as compared with reported Mg(II)adsorption at high pH conditions [13]. Based on these results andthe cost associated with the rinsing water it was decided to recom-mend process operation with 20–30 ml of rinsing water per dry gof slurry solids.

3.3. Application of the process for producing three different productMg(II) solutions

The results of the experiments aimed at achieving differentMg(II) product solutions are presented in Table 3. The experimentsin this section were carried out aiming at a 2.36 gMg/L recoverywhile using experimental conditions defined as favourable in theprevious sections (i.e. [Fe3O4] = 10 gFe/L; separation techniqueconsisting of gravity-magnetic separation followed by vacuumseparation; application of rinsing volume of 20 ml RO water/gdrained slurry followed by a second vacuum separation step(Table 1)). The high Mg/Ca ratios achieved in all the experimentsproved that precipitation of calcium-containing solids (namelyCaCO3) was almost completely avoided by the initial acidic CO2

stripping step. The Mg/Na molar ratio achieved (averaging 55 M/M) proved the high separation ability of the proposed separationmethod. With regard to the concentration of Mg(II) in the productwater, the maximal concentration in each of the attempted solu-tions is theoretically limited only by its apparent Ksp value, there-fore very high Mg(II) concentrations can be attained (for example[MgSO4] can be as high as 35.7 g/100 ml H2O [24]). The high Mg/Fe molar ratio achieved in the experiments in which CO2 was usedas the acidic substance in order to re-dissolve Mg into the productsolution are a good example that when pH is maintained above �6Fe3O4 dissolution is negligible. Operation at a high pH (but low en-ough to dissolve Mg(OH)2) is preferable and technically possiblealso when strong acids such as H2SO4 or HCl are used. In thebench-scale results presented here, pH did drop occasionally tovalues as low as �2 due to insufficiently accurate dosing controlequipment, resulting in a slightly lower Mg/Fe molar ratio. A highMg/B molar ratio was also recorded, averaging at 177. A commonimpurity in seawater-derived Mg(OH)2 precipitation processesare boron compounds [22]. Low boron content is desirable whenthe Mg(OH)2 precipitant is intended for the production of MgCl2

used in electrolysis processes to produce solid magnesium [23].Different approaches have been applied to reduce boron concen-tration in the product water, such as excessive base dosage toincrease pH to �12 and minimize B adsorption on the Mg(OH)2(s)

surface [23], rinsing of the Mg(OH)2 precipitant [22] as well asion exchange to reduce the boron content in the seawater [19].

3.4. Cost assessment

3.4.1. Assessment of operating expenses (OPEX)Estimating the operational expenses of the suggested process

comprises mainly of chemicals and energy demands, in additionto annual replenishment of the magnetite mass. With regard tochemicals demand, the strong base considered as the agent usedto precipitate Mg(OH)2 in the cost assessment calculation wasCaO, a low cost base. The value of this component in the process

0.0

0.5

1.0

1.5

2.0

2.5

3.0

0 5 10 15 20 25 30 35Mg/

Na

ratio

in p

rodu

ct s

olut

ion

(M/M

)

Fe3O4 concentration in brine (gFe/L)

Fig. 4. Mg/Na molar ratio in product solution as function of the magnetite concentration in the SWRO brine in experiments aimed at recovering 2.55 gMg/L from the brine(n = 6). Separation was effected by magnet-assisted vacuum drainage (no rinsing step applied).

O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45 43

is constant, irrespective of the type of Mg(II) solution produced. Abreakdown of the cost components in the production of CaO isshown in Table 4, based on [25]. Under the assumption that1.1 mol of CaO is used to recover 1 mol of Mg2+ (2.539 tons ofCaO per one ton of Mg(II), considering magnesium recovery effi-ciency of �90%), the cost of the base was approximated at446.6$/ton Mg(II).

Table 5 lists the estimated operational expenses for the produc-tion of 1 ton of Mg(II) by the suggested process. The required aciddosage in the process is divided into (1) Usage of strong acid in the1st step for acidification and subsequent elimination of the inor-ganic carbon concentration in the brine. This step requires�6.5 eq H+/m3 brine to reach pH 4.3 (calculated using PHREEQC[17]). Assuming that 2.4 kg Mg(II) are recovered in the processfrom each m3 of brine, the initial acidification step requires2708 eq of acid for each ton of Mg(II) recovered; (2) Either strongor weak acid may be used in the last step of the process tore-dissolve the Mg(OH)2 adsorbed on the Fe3O4 into the desiredproduct solution. Naturally, the choice of the acid affects the pro-duction costs (Table 5). The amount of acid required in this stepequals the amount of Mg(II) as Mg(OH)2 that needs to be re-dis-solved, i.e. 82,304 equivalents per ton of recovered Mg(II). It mustbe noted that in the case that Mg(HCO3)2 is produced, 50% of therequired CO2 may be taken from the CO2 released in the CaO pro-duction process, thereby cutting in half the CO2 external demand.

3.4.2. Rough assessment of capital expenses (CAPEX)The capital expenses (see Table 5) comprise mainly of stirred

reactors, magnetic separators, synthesis of the initial magnetitemass required in the process and construction costs. Magnetic

0102030405060708090

0 5 10 15Mg/

Na

ratio

in p

rodu

ct s

olut

ion

(M/M

)

Rinsing water ratio

Fig. 5. Mg/Na molar ratio in the product solution as a function of the volume of rinsing wExperiments were conducted with a 10 gFe/L magnetite concentration and a 2.36 gMg/L

separation of solids at the industrial scale is well known and fullydeveloped [26]. Various magnetic separators are commerciallyavailable for a wide range of applications [26,27]. For the currentCAPEX assessment a wet permanent magnet drum separator wasconsidered, being a frequently used wet magnetic separator[27,28]. Several considerations affect the choice of a suitable mag-netic separator. These considerations are beyond this preliminarystudy and will have to be further considered for this specific appli-cation. However, the most important parameters are particle sizedistribution, distribution of magnetic properties of the particlesthat need to be separated from one another, and the requiredthroughput of the machine [27,28]. For a 10,000 ton Mg/y plantproducing 2.4 KgMg/m3

brine, 4.17 million m3/y of brine should un-dergo treatment. Assuming that the plant operates for 300 days ayear at 24 operation cycles per day, the process requires magneticseparators at a total volume of 579 m3. The cost of two 300 m3 stir-red reactors was estimated at $200,000, assuming that the cost ofone unit is $100,000 (a very conservative figure). With regard tothe magnetic separator: a 1200 mm diameter, 3000 mm widthindustrial drum separator with 280 m3/h capacity was quoted ataround $30,000 (based on commercial quotes). Since two suchunits are required, the overall cost is estimated at $60,000. Assum-ing an operational concentration of 10 gMagnetite-Fe/Lbrine, an ini-tial magnetite mass of 0.24 ton as Fe is required, i.e. $1290 (the costof 1 ton of magnetite is estimated at $5350, under the assumptionit comprises mainly of the cost of chemicals required to produce it,according to the laboratory protocol appearing in [1]). Overall, thecapital costs amount to a total of $261,300. Increasing thisestimation to $500,000 (to consider construction and otherunaccounted-for costs), and assuming 6% interest and a 20 year

20 25 30 35 40(ml RO water/g solids)

ater (ml of RO water per 1 g of vacuumed drained Mg(OH)2 coated Fe3O4 particles)recovery.

Table 3Ratio between Mg(II) and other constituents in the product water: comparison of results obtained with different acidic substances.

Ratio between Mg(II) and other constituents in product water (M/M)

Mg/Na Mg/Ca Mg/Cl Mg/S Mg/B Mg/Fe

SWRO Brine 0.12 4.9 0.1 1.89 151.5 29,530Product solution with different acidic substances H2SO4 58.7 ± 1.9 686 ± 34 43.1 ± 3.5 – 175 ± 3.3 1132 ± 230

HCl 52.1 ± 2.1 688 ± 6.5 – 229 ± 4.7 176 ± 4.4 1262 ± 288CO2(g) 55.3 ± 3.2 603 ± 150 38.4 ± 2.9 239 ± 32 181 ± 6.1 >400,000

Table 4Specification and cost breakdown of a conventional 50 ton per day single shaft CaOproduction.

Units Per ton of CaO

CaCO3 Ton 1.786Fuel oil consumption kg 103.2Electric energy kWh 30

Costsa

CaCO3 $ 71.44Fuel Oil $ 103.2Electric energy $ 1.28Total $ 176

a Calculated according to: fuel oil 1000 US$/Ton; calcite $40/Ton; 4.26 cent$/kW h.

44 O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45

average equipment serviceable lifespan, the estimated normalizedcapital cost is $4.36 per ton Mg(II), i.e. about 0.34% of the OPEX.Therefore, even under the assumption that the capital costs are100% higher than the estimation presented in this section, sincethe effect of the CAPEX on the overall cost is almost negligible, suchan error would not change the cost picture significantly.

In summary, the overall cost (OPEX and CAPEX) calculated forthe suggested process was estimated at the range 1189–1319 $/ton Mg (Table 5) for the different Mg(II) product solutions. Thisrange seems to be cost competitive, as compared to the market costsof available magnesium chemicals, particularly for the productionof MgSO4 and Mg(HCO3)2 solutions (2516 and 3000–4000$/tonMg, respectively, [29, personal knowledge]). With regard toMg(HCO3)2 the cost is particularly attractive if the CO2 producedin the lime production process is recovered and reused in the

Table 5Production cost breakdown for the proposed process for three magnesium compounds, as

Parameter Units

Chemical demandsCaO Ton chemH2SO4

HClCO2

b

OPEXChemicalsc $/ton Mg

product)Labour, electricity, water annual magnetite replenishment and

maintenanced

Total OPEX

CAPEX $/ton Mg

Total Cost CAPEX + OPEX $/ton Mgproduct)

a �95% Product purity.b HCl is used in the CO2 stripping step. 2 mol of CO2 required per 1 mol Mg(II).c Calculated according to: H2SO4 $170/ton; HCl $200/ton; CO2 $200/ton; CaO 176$/tod Assuming �10% of total OPEX. Also included are 10% annual magnetite replenishmen

and a 10 gMagnetite-Fe/Lbrine operational concentration (0.024 ton of magnetite per year), i

dissolution step supplying about 50% of the CO2 demand, therebyfurther reducing production costs. Production of MgCl2 (market costof 1176 $/ton Mg [29]) is less straightforwardly cost competitive.

4. Summary and conclusions

� A significant improvement is presented to the separation step ofa method for extraction and reuse of Mg(II) from SWRO brines.� The method is based on precipitation and adsorption of

Mg(OH)2(s) on the surface area of Fe3O4 micro-particles whichare then separated by either gravity of magnet-assisted gravityforces.� To avoid the precipitation of CaCO3 the RO brine is first acidified

to convert all carbonate species to CO2(aq) which is thereafterstripped to the atmosphere or reused in the Mg(II) desorptionstep.� Operational conditions leading to high quality Mg(II) solution

products were found to be: Fe3O4 seed concentration at10 gFe/L, magnet-assisted gravity settling, a vacuum-assistedsolids-aqueous separation step and a rinsing step using 20 mlper g of drained solids, followed by a second vacuum-assistedsolids-aqueous separation step.� Depending on the identity of the acid used for re-dissolving the

Mg(II) three product solution can be attained (MgCl2, MgSO4

and Mg(HCO3)2), all at >97% purity.� A rough cost analysis revealed that producing MgSO4 and

Mg(HCO3)2 by the method is very attractive cost-wise, whileMgCl2 can be produced at a cost that is approximately similarto equivalent commercial products.

suming production of 10,000 ton Mg/y.

Solution produceda

MgCl2 MgSO4 Mg(HCO3)2

ical/ton product as Mg(II) 2.539 2.539 2.539– 4.168 –3.099 – 0.099– – 3.786

(II) product ($/ton 1066.4 1155.2 1182.9

118.48 128.36 131.4

1184.85(302.4)

1283.56(259.1)

1314.36(218.3)

(II) product 4.36 4.36 4.36

(II) product ($/ton 1189.21(303.5)

1287.92 (260) 1318.72 (219)

n.t at 5350 $/tonMagnetite, based on a 2.4 gMg/Lbrine production, 580 m3 brine per cycle.e. 0.0129$/ton Mg(II).

O. Lehmann et al. / Chemical Engineering Journal 235 (2014) 37–45 45

Acknowledgement

This project was funded by the Joint German-Israeli WaterTechnology Research Program (BMBF/MOST).

References

[1] D. Hyun Kim, A review of desalting process techniques and economic analysisof the recovery of salts from retentates, Desalination 270 (2011) 1–8.

[2] D. Barba, V. Brandani, G.D. Giacomo, P.U. Foscolo, Magnesium oxide productionfrom concentrated brines, Desalination 33 (1980) 241–250.

[3] I.S. Al-Mutaz, K.M. Wagialla, Production of magnesium from desalinationbrines, Resour. Conserv. Recycl. 3 (1990) 231–239.

[4] D.A. Kramer, Magnesium, its Alloys and Compounds, U.S.G.S Open-File Report,2000, pp. 01–341.

[5] L. Birnhack, O. Lahav, A new post-treatment process for attaining Ca2+, Mg2+,SO2�

4 and alkalinity criteria in desalinated water, Water Res. 41 (2007) 3989–3997.

[6] E. Drioli, E. Curcio, A. Criscuoli, G.D. Profio, Integrated system of CaCO3, NaCland MgSO4�7H2O from nanofiltration retentate, J. Memb. Sci. 239 (2004) 27–38.

[7] O. Lahav, M. Telzhensky, A. Zewuhn, Y. Gendel, J. Gerth, W. Calmano, L.Birnhack, Struvite recovery from municipal-wastewater sludge centrifugesupernatant using seawater NF concentrate as a cheap Mg(II) source, Sep. Purif.Technol. 108 (2013) 103–110.

[8] A.S. Bhatti, D. Dollimore, A. Dyer, Magnesia from seawater: a review, ClayMiner. 19 (1984) 865–875.

[9] J.A. Epstein, Utilization of the Dead Sea minerals (a review), Hydrometallurgy 2(1976) 1–10.

[10] T.G. Tripp, Production of Magnesium from Great Salt Lake, Utah USA. Natl.Resour. Environ. Issues 15 (2009) 54–61 (Saline Lakes Around the World:Unique Systems with Unique Values).

[11] B.A. Bolto, T.H. Spurling, Water Purification with magnetic particles, Environ.Monit. Assess. 19 (1991) 119–143.

[12] J. Broomberg, S. Gelinas, J.A. Finch, Z. XU, Review of magnetic carriertechnologies for metal ion removal, Mag. Electr. Sep. 9 (1998) 169–188.

[13] D.R. Dixon, Interaction of alkaline earth metal ions with magnetite, ColloidsSurf. 13 (1985) 273–286.

[14] O. Lahav, B.E. Morgan, G.R. Hearne, R.E. Loewenthal, One-step ambienttemperature ferrite process for treatment of acid mine drainage waters, J.Environ. Eng. 129 (2) (2003) 155–161.

[15] L. Petrick, Y. Dubowski, S. Klas, O. Lahav, Stable incorporation of Co2+ intoferrite structure at ambient temperature: effect of operational parameters,Water Air Soil Pollut. 190 (2008) 245–257.

[16] S. Klas, Y. Dubowski, G. Pritosiwi, J. Gerth, W. Calmano, O. Lahav, Extent andmechanism of metal ion incorporation into precipitated ferrites, J. ColloidInterface Sci. 358 (2011). pp. 129–135.

[17] D.L. Parkhurst, C.A.J. Appelo, User’s guide to PHREEQC (version 2) – A computerprogram for speciation, batch-reaction, one-dimensional transport, andinverse geochemical calculations: U.S. Geological Survey Water-ResourcesInvestigations Report, 1990, pp. 99–4259, 312.

[18] APHA, Standard Methods for the Examination of Water and Wastewater, 20thEd. (Method 2710D). American Public Health Association, American WaterWorks Association and Water Environmental Federation, Washington, DC,USA, 2011.

[19] M. Seeger, W. Otto, W. Flick, F. Bickelhaupt, O.S. Akkerman, Magnesiumcompounds, Ullmann’s Encycl. Ind., Chem. 22 (2011) 41–74.

[20] S. Klas, Y. Dubowski, O. Lahav, Chemical stability and extent of isomorphoussubstitution in ferrites precipitated under ambient temperatures, J. HazardMater. 193 (2011) 59–64.

[21] B.E. Morgan, O. Lahav, R.E. Loewenthal, Advances in seeded ambienttemperature ferrite formation for treatment of acid mine drainage, Environ.Sci. Technol. 39 (19) (2005) 7678–7683.

[22] V. Martinac, M. Labor, N. Patric, Boric oxide is seawater derived magnesia,Indian J. Geomarine Sci. 33 (3) (2004) 226–230.

[23] H.A. Robinson, R.E. Friedrich, R.S. Spencer, Dow chemical Company,Magnesium Hydroxide from sea water, US Patent Office, 1943, Serialnumber 492,860.

[24] M. Haynes (Ed.), CRC Handbook of Chemistry and Physics, 93rd Edition(Internet Version 2013), CRC Press/Taylor and Francis, Boca Raton, FL.

[25] A. Meier, N. Gremaud, A. Steinfeld, Economic evaluation of the industrial solarproduction of lime, Energy Conver. Manage. 46 (2005) 905–926.

[26] C.T. Yavuz, A. Prakash, J.T. Mayo, V.L. Colvin, Magnetic separations: from steelplants to biotechnology, Chem. Eng. Sci. 64 (2009) 2510–2521.

[27] J. Svoboda, T. Fujita, Recent developments in magnetic methods of materialSeparation, Miner. Eng. 16 (9) (2003) 785–792.

[28] M. Dworzanowski, Optimizing the performance of wet drum magneticseparators, J. Southern African Inst. Min. Metall. 110 (11) (2010) 643–653.

[29] M. Telzhensky, L. Birnhack, O. Lehmann, E. Windler, O. Lahav, Selectiveseparation of seawater Mg2+ ions for use in downstream water treatmentprocesses, Chem. Eng. J. 175 (2011) 136–143.