Electrochemical Induced Calcium Phosphate Precipitation: Importance

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Electrochemical induced calcium phosphate precipitation: importance of local pH Yang Lei, Bingnan Song, Renata van der Weijden, Michel Saakes, and Cees J.N. Buisman Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b03909 • Publication Date (Web): 05 Sep 2017 Downloaded from http://pubs.acs.org on September 6, 2017

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Environmental Science & Technology

1

Electrochemical induced calcium phosphate precipitation: importance of local

2

pH

3

Yang Lei†, ‡, Bingnan Song†, ‡, Renata D. van der Weijden*, †, ‡, Michel Saakes†, Cees

4

J.N. Buisman†, ‡

5

†

6

8900CC Leeuwarden, The Netherlands

7

‡

8

P.O. Box 17, 6700AA Wageningen, The Netherlands

9

*

Wetsus, Centre of Excellence for Sustainable Water Technology, P.O. Box 1113,

Sub-department Environmental Technology, Wageningen University and Research,

Corresponding author

10

Email: [email protected]

11

ABSTRACT

12

Phosphorus (P) is an essential nutrient for living organisms and cannot be replaced or

13

substituted. In this paper, we present a simple yet efficient membrane free

14

electrochemical system for P removal and recovery as calcium phosphate (CaP). This

15

method relies on in-situ formation of hydroxide ions by electro mediated water

16

reduction at a titanium cathode surface. The in-situ raised pH at the cathode provides

17

a local environment where CaP will become highly supersaturated. Therefore,

18

homogeneous and heterogeneous nucleation of CaP occurs near and at the cathode

19

surface. Because of the local high pH, the P removal behavior is not sensitive to bulk

20

solution pH and therefore, efficient P removal was observed in three studied bulk

21

solutions with pH of 4.0 (56.1%), 8.2 (57.4%) and 10.0 (48.4%) after 24 hours of

22

reaction time. While P removal efficiencies are not generally affected by bulk solution 1

ACS Paragon Plus Environment


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pH, the chemical-physical properties of CaP solids collected on the cathode are still

24

related to bulk solution pH, as confirmed by structure characterizations. High initial

25

solution pH promotes the formation of more crystalline products with relatively high

26

Ca/P ratio. The Ca/P ratio increases from 1.30 (pH 4.0) to 1.38 (pH 8.2) and further

27

increases to 1.55 (pH 10.0). The formation of CaP precipitates was a typical

28

crystallization process, with an amorphous phase formed at the initial stage which

29

then transforms to the most stable crystal phase, hydroxyapatite, which is inferred

30

from the increased Ca/P ratio from 1.38 (day 1) to the theoretical 1.76 (day 11) and by

31

the formation of needle-like crystals. Finally, we demonstrated the efficiency of this

32

system for real wastewater. This, together with the fact that the electrochemical

33

method can work at low bulk pH, without dosing chemicals and a need for a

34

separation process, highlights the potential application of the electrochemical method

35

for P removal and recovery.

36

INTRODUCTION

37

Phosphorus (P) is an irreplaceable nutrient, but it is also associated with

38

eutrophication.1-4 Indeed, on the one hand, a large amount of P is discharged to

39

surface waters resulting in eutrophication due to limited recycling.3 On the other hand,

40

the quantity and the quality of P ore has declined in the past decades because of P

41

rock mining for producing fertilizers.1 Evelyn Desmid’s calculation which applied the

42

data of U.S geological survey 2012, suggests that natural P reserves will be fully

43

depleted in 372 years if current mining rates are maintained.1 In addition, considering

44

the uneven distribution of P rock reserves, there may arise a P shortage for countries 2


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that completely depend on importing P rock in the near future.4, 5 The potential P

46

shortage along with P discharge associated eutrophication, has created increased

47

awareness of the importance of P recycling.1,

48

government has set a national goal to recover at least 40% of P in wastewater

49

treatment plants.7

50

There are many P removal methods available,1, 8-10 but efficient and economically

51

feasible P recovery methods are quite limited. Among the few methods, struvite

52

(MgNH4PO4·6H2O) formation and precipitation is regarded as one of the most

53

promising ways.6,

54

bioavailability than iron and aluminum phosphate.10 However, it is necessary to

55

supply a Mg source to assist struvite formation,13-15 which makes the struvite process

56

less economically attractive because of the low concentration of Mg2+ in wastewater.15

57

Alternatively, calcium phosphate (CaP), which is the mined component in P rock,

58

would be a better solution.16, 17 CaP solids can form without adding Ca2+ since there is

59

often already sufficient Ca2+ in bodies of water.18 Therefore, P recovery via CaP

60

formation and precipitation is a preferred method, and has received a lot of

61

attention.17, 19

62

CaP precipitation is a very complex process. In general, the process is controlled by

63

the chemical species in solution, including Ca and P concentrations and pH.

64

induce CaP precipitation, the solution needs to be highly supersaturated. The typical

65

way to create a supersaturated condition is by adding caustic soda to increase solution

66

pH. However, because wastewater normally has a considerable buffering capacity

11-13

4, 6

For instance, the Swedish

Struvite, which is a slow-release fertilizer, shows higher

20-22

To

3



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67

because of the presence of organic acids and inorganic carbonates, significant base

68

addition is needed in order to increase the bulk solution pH to a certain level that

69

would induce CaP precipitation. For instance, as reported by Jaffer et al.,23 the sodium

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hydroxide addition is accounted for up to 97% of the total chemical costs associated

71

with P recovery by struvite formation method. Furthermore, the traditional chemical

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precipitation based methods produce a large quantity of sludge, which still needs to be

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treated before recycling.24

74

Recently, (bio)electrochemical processes were suggested as next generation

75

technologies for treating (in)organic polluted water

76

strategy for nutrient removal and recovery from nutrient rich wastewater.26 Though

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(bio)electrochemical reactions are quite complicated processes, they can be simply

78

divided as anode oxidation and cathode reduction. Most environment related

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electrochemical applications depend on the processes at the anode. The

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well-established electro-Fenton method for degrading organic pollutants is a good

81

example.27 By contrast, the role of cathode mediated reduction, has just begun to be

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explored for remediation and recovery by environmental scientists.28 The

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(bio)electrochemical induced P removal and recovery as struvite has been

84

well-documented.29-31 However, the electrochemical assisted struvite formation, like

85

chemical precipitation still relies on dosing of costly Mg2+. Moreover, the importance

86

of local pH at the cathode with respect to electrochemical P recovery has not been

87

recognized yet. Most studies mention that the increased pH is responsible for the

88

precipitation of phosphate salts but none, to the best of our knowledge, has

25

and recognized as an efficient

4


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89

investigated the role of local pH in detail. This is because it is difficult to measure the

90

local pH directly, as there still are no reliable pH sensors for detecting the

91

electrochemically induced local pH at the electrode surface, though there are some

92

special designed lab tools.32, 33 Moreover, the importance of local pH was seemingly

93

ignored. Some researchers equate bulk solution pH to local pH and therefore just

94

record bulk solution pH and use it as the pH for phosphate salts precipitation.29, 34

95

Consequently, the reported results with respect to local pH varied from experiment to

96

experiment. As an example, Wang et al.29 reported the slight increase of pH near the

97

cathode from 7.0 to 7.5 as the cause for pure struvite formation in their

98

electrochemical system. However, the local pH can be much higher than can

99

measured.33

be

100

To the best of our knowledge the electrochemical induced CaP precipitation on the

101

cathode has not been reported, in terms of P removal and recovery and at various bulk

102

pH. Although CaP coverage of the cathode might seem unwanted, we see this as an

103

opportunity to separate P from waste streams with low P concentrations. Therefore,

104

the purpose of this study, was to evaluate the efficiency of a single electrochemical

105

cell without membrane for P removal and recovery by forming CaP precipitates. The

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importance of a local high pH in the electrochemical cell was demonstrated by

107

evaluating the performance of this system at low, higher and high bulk solution pH

108

combined with theoretical calculations. Finally, the efficiency and the cost for

109

treatment of real wastewater were addressed to evaluate the potential for this new P

110

recovery system. 5



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MATERIALS AND METHODS

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Materials. All chemicals used here were at least reagent-grade. Di-sodium

113

monohydrogen phosphate (Na2HPO4) and sodium sulphate anhydrous (Na2SO4) were

114

purchased

115

(Ca(NO3)2·4H2O) was received from Merck (Germany). Electrodes were provided by

116

MAGNETO Special Anodes BV (Schiedam, The Netherlands).

117

Electrolysis Setup. The electrochemical cell consisted of two compartments, one

118

working cell (500 mL) for CaP precipitation and one tank cell (500 mL) for mixing

119

and sampling. The total solution in the two compartments (1000 mL) is circulated

120

with a pump at a flow rate of 100 mL/min. The anode material is platinum coated (50

121

g/m2) titanium mesh with a round shape (Ø 10 cm, thickness 0.1cm) and it is

122

perpendicularly welded to a 10 cm long Ti rod (Ø 0.3 cm). The cathode is a pure

123

titanium plate similarly welded (grade A, Ø 8.2 cm, thickness 0.1 cm). A pH sensor

124

was placed in the sampling tank to record bulk solution pH change. In some cases, the

125

pH electrode (Ø 1.2 cm, Endress Hauser, Germany) was also placed near the cathode

126

(~ 1.0 mm), in order to record local pH. The pH sensors were calibrated weekly. The

127

diagram of the set sup is shown in Figure S1.

128

Electrolysis Experiments. The electrochemical precipitation process was conducted

129

with 0.6 mM P and 1.0 mM Ca under constant current (20 mA) conditions and at

130

constant ionic strength mediated by 50 mM Na2SO4. The choice for a sulfate salt was

131

made because it does not interfere with the precipitation of CaP and does not produce

132

harmful chlorine gas as well. While the initial Ca concentration is close to its natural

from

VWR

(Leuven,

Belgium).

Calcium

nitrate

tetrahydrate

6


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133

concentration, the initial P concentration was higher compared to real wastewater in

134

order to collect sufficient solid samples for further characterization. Where

135

appropriate, the bulk solution pH was adjusted by concentrated NaOH or HNO3.

136

Unless specified, the electrolysis process was open to air and lasted for 24 hours at

137

room temperature. The bulk solution pH was monitored during the whole process and

138

logged by a computer program (Liquisys M CPM 253, Endress + Hauser, Naarden,

139

The Netherlands).

140

Calcium phosphate collection. After the reaction was stopped at a predetermined

141

time, the solutions in the electro cell were carefully removed with a syringe as to not

142

disturb CaP precipitates at the cathode surface, for the sake of solid characterization.

143

Then the electrode with precipitates on its surface was air dried. After drying, CaP

144

solids were harvested by light scraping. After sampling, the cathode was immersed in

145

a 1.0 M HNO3 solution to remove any CaP remaining and then thoroughly rinsed with

146

Milli-Q water and dried again for use.

147

Analytical methods. The concentrations of P and Ca ions were analyzed by ICP-AES.

148

X-ray diffraction (XRD) was used to identify the crystal structure (or absence thereof

149

if amorphous) and collected on a Bruker D8 advanced diffractometer equipped with a

150

Vantec position sensitive detector and with a Co Kα radiation (λ = 0.179 nm) over a

151

range of 5−90° in 0.02 step sizes with an integration time of 0.5 s. Raman spectra

152

were obtained using a LabRAM HR Raman spectrometer from Horiba Jobin Yvon to

153

obtain bonding information of collected solids. This system is equipped with a

154

mpc3000 laser emitting at 532.2 nm and an 800 mm focal length achromatic flat field 7



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monochromator (grating of 600 grooves·mm-1). The laser beam was focused on the

156

sample with an Olympus Bx41 microscope equipped with a 50x objective lens, which

157

gives a spot size ca.1-2 µm and resolution of 6 cm-1. The detector is a Synapse

158

multichannel air cooled (-70°C) CCD. The applied laser power was between 5 and 50

159

mW (using density filters). The measurement time varied 5 to 30 s. Finally, the data

160

were processed with LabSpec software. The morphology of collected products and

161

their elemental compositions were examined by a Scanning Electron Microscope

162

(SEM) coupled with Energy dispersive x-ray spectroscopy (EDS) (JEOL-6480LV,

163

JEOL Ltd., Japan). Samples were coated with gold using a JEOL JFC-1200 Fine

164

coater at 10 Pa for 30 s.

165

Calculations. The degree of saturation (Ω) and saturation index (SI) of a solution

166

regarding a mineral phase, are defined as follows:35

167

Ω =

168

SI = log ()

169

Where IAP refers to the ion activity of the associated lattice ions and Ksp is the

170

thermodynamic solubility product. The computer program visual MINTEQ36 was

171

applied to calculate SI, as an indication for the potential saturation of possible

172

products. Ca and P fractions were acquired by using Hydra–Medusa database.37

173

Based on Faraday’s law of electrolysis assuming that the electricity consumed was

174

100% used for water reduction and meanwhile supposing the produced OH− was not

175

consumed by other occurring reactions and was homogenously mixed in the local

176

layer, the theoretical maximum local pH, with respect to the thickness of local layer (δ,

(1)

(2)

8


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m) and electrolysis time (t, s) can be calculated by equations 3-5:

178

( ) =

179

[ ] =

180

# = 14 + log (

181

I electricity current (A); z number of electrons transferred in the solution, z = 1; F

182

Faraday constant 96,485 (C/mol); d diameter of cathode (d = 0.082 m). It should be

183

noted here that the real local pH will be below the theoretical calculated value because

184

the current efficiency is unlikely to reach 100% and the electrochemically produced

185

H+ at anode will react with OH− to a certain extent.

186

RESULTS AND DISCUSSION

187

Effects of initial bulk pH (pH0). As a proof of principle, recovery of P in the

188

electrochemical system was evaluated at three pH values including background

189

solution pH (pH0 ~ 8.2) after mixing of all chemicals, weak acidic (pH0 4.0) and

190

alkaline (pH0 10.0) conditions. As can be seen from Figure 1A, under open circuit

191

conditions, only 20% of P was removed in the case of pH0 10.0 and there was no

192

obvious P removal at pH0 4.0 and 8.2. For pH0 4.0, the solution was undersaturated

193

with respect to hydroxyapatite (HAP) (SIHAP = −15.5) and with respect to any calcium

194

solid species like gypsum (Figure 1B). In addition, the calculation of the species

195

distribution showed that nearly 87% of Ca existed as dissolved CaSO4 and P was

196

almost 100% present as H2PO4− (Figure S2). Therefore, it is not surprising that no

197

CaP precipitated from solution at pH0 4.0. In terms of pH0 8.2, while the

198

thermodynamic calculation indicates the solution is supersaturated with respect to

(3)

! "

(4) '

)

!"

(5)

9



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199

HAP (SIHAP = 8.6) and the fraction calculation also suggests the formation of HAP

200

(Figure S2), no visible precipitates were found in reactors. Actually, many lakes are

201

also supersaturated with respect to HAP without HAP being found in the lake

202

sediments.35 Indeed, thermodynamic predictions for precipitation of certain solids do

203

not imply that they are kinetically favorable. The precipitation rate may be too slow to

204

be observed and precipitation may progress via the Ostwald Step Rule. Interestingly,

205

it was found that the application of a low current (20 mA, current density corresponds

206

to 3.79 A/m2) makes a big difference for removal of P. The P removal efficiencies

207

jumped to over 48% in all cases; 56.1%, 57.4% and 48.4% of P was removed at pH0

208

4.0, pH0 8.2 and pH0 10.0 respectively (Figure 1A) within 24 hours. We found that

209

approximately 50% of Ca was removed as well. The simultaneous removal behavior

210

of Ca and P indicates the removal as CaP precipitates. The precipitated solids were

211

characterized with XRD and Raman spectroscopy. The Raman spectrum (Figure 1C)

212

of the three samples almost all show internal bands of CaP, including a main ν1 PO43−

213

peak around 955 cm−1 and well isolated ν2 PO43− (~425 cm−1) and ν4 PO43− (~590

214

cm−1) peaks, which clearly demonstrates the formation of CaP.38, 39 Interestingly, the

215

XRD patterns (Figure 1D) suggest amorphous products are produced in acid and

216

neutral solution as confirmed by the lack of sharp peaks and the presence of a broad

217

peak around 38° though at pH0 10.0, a relatively more crystalline product is formed.

218

The sharp peak around 30° indicates the presence of more crystalline CaP phases.

219

However, the product is still dominantly amorphous. Most of the sharp peaks of pH0

220

10.0, unfortunately, is attributed to Na2SO4 because the electrode was air-dried 10


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221

without rinsing.

222

While it is not surprising that P was removed in an alkaline solution, the high removal

223

efficiency of P at pH0 4.0 was not expected. As seen from Figure 1B, the solution at

224

pH0 4.0 is undersaturated for all possible CaP products. The only factor that can

225

contribute to the increase of SI here could be the increase of pH. By contrast, Figure

226

1E shows the solution pH decreases largely for pH0 8.2 and pH0 10.0, in which the

227

solution pH drops to 4.6 and 4.0 respectively. Regarding pH0 4.0, it also declines to

228

3.4 after 24 hours’ reaction. It should be noted here that under open circuit the

229

solution pHs also drop to some extent due to equilibration with atmospheric CO2 in an

230

open system (Figure 1E). In conclusion, it may be reasonable to infer that bulk

231

solution pH is not that important, in terms of P removal efficiency.

232

Importance of local pH. A phenomenon that we observed during our experiments is

233

that precipitates just formed at/near the surface of the cathode. This points to different

234

conditions at the cathode surface than in the bulk solution. The possible differences

235

could be pH, Ca and P concentration, which determines the saturation of CaP species

236

in our system. Indeed, electro migration could transfer negative ions to the anode and

237

positive ions to cathode. However, because the relative low concentration of Ca2+

238

compared to electrolyte (50 mM Na2SO4), it is unlikely that Ca2+ can be enriched to

239

such extent that it can increase the saturation state of CaP. Also, if electro migration

240

plays an important role here, the P concentration in the vicinity of cathode surface

241

should decline correspondingly. Therefore, it is concluded that electro migration of

242

ions does not play a crucial role in this system. The only possible reason for 11



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243

precipitation should then be attributed to the production of OH− by electrochemical

244

mediated water reduction at the cathode:

245

2H2O + 2e− → 2OH− + H2↑

(6)

246

Though the produced OH− will diffuse to the bulk solution and the diffusion rate will

247

increase with mixing rate, the relatively high pH in the very vicinity of cathode will

248

not disappear.40 While we did not have special pH sensors to record local pH, an

249

attempt was made to measure the local pH by a general pH sensor. Indeed, a big

250

difference was found between bulk solution pH and the so measured local pH, as

251

shown in Figure 2A. For example, in 1 hour, while the solution pH dropped to 7.4

252

from 8.2, the local pH went up to 9.9. However, as the measurement of local pH by

253

this method is sensitive to the distance between the sensor and cathode, it is difficult

254

to record a consistent pH. Consequently, the trend of local pH changes a lot. Indeed,

255

though we did not measure the exact thickness of the precipitation layer, it is

256

supposed that the local crystallization zone ranges to less than 1 mm away from the

257

cathode surface, which was proven by a simple test. When we put a glass plate (26 ×

258

26 × 1mm) on the cathode surface, covering 12.8% of the cathode, there was no

259

precipitates initiated from the glass surface. This showed that CaP precipitation just

260

take places in the local region of the Ti cathode where the surface pH is much higher

261

than the bulk solution pH because of the electrochemical production of hydroxide ions.

262

Considering the size of a regular pH sensor as used in our experiments and the

263

thickness of the reaction zone where a high local pH is created, it is evident that the

264

local pH cannot be recorded consistently with a common pH electrode. Nevertheless, 12


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there is no doubt that local pH is much higher than bulk solution pH. In addition to

266

measuring the local pH directly, an attempt was made to calculate the local pH

267

theoretically. The production of OH− corresponds to the electricity consumed with

268

time elapse and can be calculated by Faraday’s law. The calculation results (Figure 2B)

269

suggest that the local pH decreases with the thickness of local diffusion layer and it

270

can reach pH values as high as 13.2 and 14.5 theoretically for an assumed maximum

271

thickness of the local diffusion layer of 1 mm and after, respectively, 1 and 24 hours

272

electrolysis. The local pH can be even higher if we assume a smaller local diffusion

273

layer. The theoretical calculation along with the fact that CaP only forms in the

274

vicinity of and on the cathode surface indicated that the electrochemically induced

275

high local pH is indeed responsible for the phosphate precipitation.

276

Crystallization mechanism. As discussed above, the bulk solution pH values in the

277

electrochemical system are not as important as in traditional chemical precipitation

278

processes. This is attributed to the electrochemically created difference between bulk

279

solution pH and the local pH at the vicinity of cathode. A possible CaP formation and

280

precipitation mechanism based on the increase of local pH is suggested here. For the

281

first step, the consumption of electrons by cathode mediated water reduction, created

282

the high local pH (see eq 6). Meanwhile, dihydrogen phosphate (H2PO4−) reacts to

283

monohydrogen phosphate (HPO42−) and phosphate (PO43−) via acid-base reactions in

284

the local area.

285

xCa2+ + yHPO42−/PO43− + nH2O → ACP↓

(7)

286

Ca2+ + HPO42− + 2H2O → DCPD↓

(8) 13



287

8Ca2+ + 4PO43− + 2HPO42− + 5H2O → OCP↓

288

10Ca2+ + 6PO43− + 2OH− → HAP↓

Page 14 of 28

(9) (10)

289

OCP/DCPD/ACP + xCa2+ + yOH− → HAP↓

290

In the second step, homogenous nucleation of CaP occurs because of the increased

291

thermodynamic driving force and the declined solubility of CaP salts, both resulting

292

from the high local pH. It should be noted that the Ti cathode might also provide a

293

favorable surface for CaP nucleation in this system. Even so, it takes more than four

294

hours to see macroscopic precipitates. These then promote the growth and

295

precipitation of precursor phases of HAP. The formed precipitates were weakly

296

attached to the cathode surface via electrostatic interactions and continued to

297

growing.41 Gradually, the precipitates covered the cathode surface. One may worry

298

that covering the cathode surface with CaP precipitates will increase the resistance

299

and will corrupt the local pH and thus under constant current conditions, the cell

300

potential would increase a lot. However, this phenomenon was not observed in our

301

system, probably because the surface is not completely blocked as a result of the

302

formation of hydrogen bubbles that resulted in small channels through the CaP layer.

303

In addition, because of the design of our electrodes, the bottom side (or even the rod)

304

of the cathode can work equally well when the top of the cathode is covered.

305

The possible intermediate phases, including amorphous calcium phosphate (ACP),

306

brushite

307

(Ca8(HPO4)2(PO4)4·5H2O, OCP) can be involved in the crystallization process (see

308

eqns. 7 to 9). However, we were not able to characterize all possible species

(CaHPO4·2H2O,

DCPD),

(11)

and

octacalcium

phosphate

14


Page 15 of 28


309

mentioned. The associated initial phase in our system was demonstrated as ACP by

310

the absence of peaks in the corresponding XRD patterns. The typical broad peak at 2θ

311

= 38° confirms the formation of ACP as a precursor (Figure 3A). Regarding ACP,

312

there is no defined chemical formula yet but normally the formula of Ca9(PO4)6·nH2O

313

is used since Posner and Betts proposed that structure.42 However, the Ca/P ratio (1.38,

314

Figure S3) in our system is lower than the proposed value and therefore, the formula

315

of CaxHy(PO4)z·nH2O is suggested. The formation of ACP in our system can be

316

expressed as given in eq 7. In addition, carbonate, which could originate from

317

atmospheric CO2 under alkaline conditions, might also be incorporated or precipitate

318

as calcium carbonate. However, both XRD and Raman data cannot confirm the

319

presence of CaCO3. The formation of ACP in our system agrees with Ostwald rule,43

320

which foresees that the crystallization process is initiated by the formation of least

321

thermodynamically stable phase. Indeed, though thermodynamics predict HAP

322

formation, the direct formation of HAP was not observed. This is because the

323

formation of HAP is much slower than that of either ACP or OCP, and during

324

simultaneous phase formation, a larger percentage of the kinetically favored species

325

may be observed, even though it has a much smaller thermodynamic driving force.44

326

At constant temperature, the transformation kinetics is a function of only pH because

327

pH regulates both the dissolution of precursor phases and the formation of the early

328

HAP nuclei.44 In our system, the cathode mediated water reduction regulates the

329

production of OH−. Therefore, we thought that when the electrolysis time is increased,

330

the initially formed ACP and other intermediate CaP phases may transform to HAP 15



Page 16 of 28

331

via eqns. 9-11.

332

To check if HAP can form eventually, we increased the electrolysis time up to 11 days.

333

The results, however, illustrate that even after a period of 4 days, the products were

334

still dominantly amorphous (Figure 3A). This indicates that the precursor phase can

335

be maintained for a long period. However, we found that though the phase does not

336

change, the solids particle size increased, as can be seen from SEM images shown in

337

Figure 3B. Note that these SEM images have the same magnification factor (× 10000).

338

In addition to the growth of particles, the corresponding Ca/P ratio also increases to

339

1.44 (See Figure S3). However, on day 7, both the morphology and phase changed.

340

The XRD data (Figure 3A) along with the typical needle-like shape45 (Figure 3B)

341

demonstrates the formation of HAP on the 7th day. The good match with peaks around

342

13°, 30° and the four peaks in the range of 2θ 38° to 42° for commercial HAP

343

confirms the transformation to HAP. The Ca/P ratio (1.66) on day 7 also agrees well

344

with theoretical Ca/P ratio (1.67). On day 11, the particle size increased again and the

345

Ca/P ratio reached 1.76, but the morphology remained need-like. The phase

346

transformation to HAP can be further supported by Raman data (Figure 3C), where

347

the ν1 PO43− band shifted from 955 cm−1 typical for ACP (day 1 and 4) to 963 cm−1

348

that is for HAP (day 7 and 11).46 In addition to solid characterization and analysis, the

349

changes of Ca and P concentrations in the bulk solution also support the phase

350

transformation. Figure 3C shows the removal trend of Ca and P from solution. Both P

351

and Ca concentrations decreased with electrolysis time. After 7 days, more than 90%

352

P and Ca precipitated from solution. Specifically, at the end of all reaction periods, the 16


Page 17 of 28


353

removal efficiency of P is higher than of Ca, but the difference for 7 days and 11 days

354

(3.1% ± 0.3) is much lower than for day 1-7 (9.5% ± 1.0). This result suggests that

355

low Ca/P ratio products (ACP) are formed initially on day 1 (Ca/P = 1.38) and day 4

356

(Ca/P = 1.44) and later transformed into high ratio (1.66 and 1.76 for day 7 and 11

357

respectively) product (HAP), thanks to the continuous production of OH− at the

358

cathode surface. Because the initial molar ratio of Ca (1.0 mM) to P (0.6 mM) is 1.67

359

and therefore the formation of low ratio Ca/P products will result in the relatively

360

lower use of Ca. To conclude, the formation of HAP in the electrochemical system is

361

identified as a typical crystallization process, starting with an amorphous phase

362

followed by the precursors and finally transformed to the thermodynamically most

363

stable phase (HAP).

364

Electrochemical recovery of phosphorus in real wastewater. Besides studying the

365

efficiency and the precipitation mechanism using simulated solutions with various

366

bulk pH, the efficiency of electrochemical P precipitation for real wastewater was

367

investigated and compared with conventional chemical precipitation, in terms of

368

efficiency and cost. Detailed information about the wastewater compositions,

369

experimental methods and cost calculation are provided in SI (See the texts and Table

370

S1).

371

Figure 4 gives a summary of the results of electrochemical and chemical precipitation.

372

In electrochemical precipitation system, after a period of 24, 48 and 72 hours, the P

373

concentration decreased from 8.0 to 4.3, 3.1 and 2.3 mg/L respectively. This

374

corresponds to a removal efficiency of 42.8% in 24 h, 62.1% in 48 h and 71.5% in 72 17



Page 18 of 28

375

h. Though the wastewater has a complicated matrix (see Table S1) and a much lower

376

P concentration, the removal efficiency is comparable to the simulated solutions. This

377

is probably due to the role of Mg and Ca. In the wastewater, the removal of P results

378

from both calcium phosphate and magnesium phosphate precipitation. This was

379

concluded from the simultaneous removal of P, Ca and Mg (Figure 4A). At the same

380

time, we found the concentration of inorganic carbon also decreased from 166 to 115

381

mg/L (Figure S4). This points to formation and precipitation of CaCO3 and MgCO3 or

382

a mixed phase. The contribution of CaCO3 was also reported on P removal from

383

wastewater by Ca-P precipitation.16, 34 In addition to the coprecipitation of carbonate

384

salts, the heavy metal ions in the wastewater, which we did not address in this paper,

385

might be removed via adsorption or incorporation, as reported in a previous study on

386

struvite formation from urine.47 Hence, for the purpose of P recycling for use in

387

fertilizer, the behavior of toxic ions in the phosphate recovery process should be

388

investigated in detail. Ideally, heavy metal ions (i.e., Zn, Cu) could be incorporated

389

and work as micro nutrients, but their contents should be kept below the standard for

390

P fertilizers. A more in depth study on the fate and behavior of coexisting components

391

and the corresponding effects on the possible application of products is ongoing.

392

Electrochemical P precipitation was also compared with conventional chemical

393

precipitation, in terms of efficiency and cost. Clearly, as shown in Figure 4, as

394

expected, chemical precipitation is more efficient than electrochemical precipitation

395

regarding removal efficiency. After adjusting the solution pH ≥10, over 78.8% P

396

(Figure 4B) was removed from the solution. It should be noted that the P removal 18


Page 19 of 28


397

refers to the P removal after filtration through 0.45 µm membrane and therefore this

398

value is higher than the precipitation efficiency (see Figure 4B), as the formed

399

products do not have a good settling rate. For example, the removal efficiency of P is

400

93.9% at pH 11 but the corresponding precipitation efficiency is only 67.8%. Hence,

401

in chemical precipitation process, a follow up separation process is needed. However,

402

in the electrochemical system, because the precipitation only happens near and on the

403

cathode surface, removal and separation are simultaneous. The extra separation

404

process is therefore avoided.

405

For cost comparison, we only considered the electricity cost in the electrochemical

406

system and the caustic soda cost for the chemical precipitation system. After

407

normalizing the cost as €/kg P, the cost of electrochemical precipitation is 41 €/kg P,

408

which is comparable to chemical precipitation. The cost of chemical precipitation

409

depends on the solution pH and varies from 18.9 to 61.1 €/kg P. The lowest cost is

410

achieved at pH 10. As the cost of the two methods are of the same magnitude, we

411

believe optimization of the electrochemical process can make the process

412

economically viable.

413

ASSOCIATED CONTENT

414

This material is available free of charge on ACS publication website

415

(http://pubs.acs.org/).

416

ACKNOWLEDGMENTS

417

This work was performed in the cooperation framework of Wetsus, European Centre

418

of Excellence for Sustainable Water Technology (www.wetsus.eu). Wetsus is 19



419

co-funded by the Dutch Ministry of Economic Affairs and Ministry of Infrastructure

420

and Environment, the European Union Regional Development Fund, the Province of

421

Fryslân, and the Northern Netherlands Provinces. This research has received funding

422

from the European Union’s Horizon 2020 research and innovation programme under

423

the Marie Skłodowska-Curie grant agreement No 665874. The authors are grateful to

424

the participants of the research theme “Resource Recovery” for fruitful discussions

425

and financial support.

426

REFERENCES

427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455

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27. Brillas, E.; Sirés, I.; Oturan, M. A., Electro-Fenton process and related electrochemical technologies based on Fenton’s reaction chemistry. Chem Rev 2009, 109, (12), 6570-6631. 28. Heijne, A. T.; Liu, F.; Weijden, R. v. d.; Weijma, J.; Buisman, C. J.; Hamelers, H. V., Copper recovery combined with electricity production in a microbial fuel cell. Environ Sci Tech 2010, 44, (11), 4376-4381. 29. Wang, C. C.; Hao, X. D.; Guo, G. S.; van Loosdrecht, M. C. M., Formation of pure struvite at neutral pH by electrochemical deposition. Chem Eng J 2010, 159, (1-3), 280-283. 30. Cusick, R. D.; Logan, B. E., Phosphate recovery as struvite within a single chamber microbial electrolysis cell. Bioresour Technol 2012, 107, 110-115. 31. Cusick, R. D.; Ullery, M. L.; Dempsey, B. A.; Logan, B. E., Electrochemical struvite precipitation from digestate with a fluidized bed cathode microbial electrolysis cell. Water Res 2014, 54, 297-306. 32. Zhang, J.; Lin, C.; Feng, Z.; Tian, Z., Mechanistic studies of electrodeposition for bioceramic coatings of calcium phosphates by an in situ pH-microsensor technique. J Electroanal Chem 1998, 452, (2), 235-240. 33. Honda, T.; Murase, K.; Hirato, T.; Awakura, Y., pH measurement in the vicinity of a cathode evolving hydrogen gas using an antimony microelectrode. J Appl Electrochem 1998, 28, (6), 617-622. 34. Kappel, C.; Yasadi, K.; Temmink, H.; Metz, S. J.; Kemperman, A. J. B.; Nijmeijer, K.; Zwijnenburg, A.; Witkamp, G. J.; Rijnaarts, H. H. M., Electrochemical phosphate recovery from nanofiltration concentrates. Sep Purif Technol 2013, 120, 437-444. 35. House, W., The physico-chemical conditions for the precipitation of phosphate with calcium. Environ Technol 1999, 20, (7), 727-733. 36. Gustafsson, J., Visual MINTEQ ver. 3.0. Department of Land and Water Resources Engineering, Royal Institute of Technology: Stokholm, Sweden 2011. 37. Puigdomènech, I., Chemical equilibrium software Hydra and Medusa. Inorganic Chemistry Department, Technology Institute, Stockholm, Sweden 2001. 38. Akiva, A.; Kerschnitzki, M.; Pinkas, I.; Wagermaier, W.; Yaniv, K.; Fratzl, P.; Addadi, L.; Weiner, S., Mineral formation in the larval zebrafish tail bone occurs via an acidic disordered calcium phosphate phase. J Am Chem Soc 2016, 138, (43), 14481-14487. 39. Ensikat, H.-J.; Geisler, T.; Weigend, M., A first report of hydroxylated apatite as structural biomineral in Loasaceae–plants’ teeth against herbivores. Sci Rep 2016, 6. 40. Ter Heijne, A.; Strik, D. P.; Hamelers, H. V.; Buisman, C. J., Cathode potential and mass transfer determine performance of oxygen reducing biocathodes in microbial fuel cells. Environ Sci Tech 2010, 44, (18), 7151-6. 41. Chen, H.-T.; Wang, M.-C.; Chang, K.-M.; Wang, S.-H.; Shih, W.-J.; Li, W.-L., Phase Transformation and Morphology of Calcium Phosphate Prepared by Electrochemical Deposition Process Through Alkali Treatment and Calcination. Metall Mater Trans A 2013, 45, (4), 2260-2269. 42. Betts, F.; Posner, A., An X-ray radial distribution study of amorphous calcium phosphate. Mater Res Bull 1974, 9, (3), 353-360. 43. Ostwald, W., Studien über die Bildung und Umwandlung fester Körper. Z. phys. Chem 1897, 22, 289-330. 44. Wang, L.; Nancollas, G. H., Calcium orthophosphates: crystallization and dissolution. 22


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TOC Graphic.

24


Page 25 of 28


553

Figure 1. 30

80

60 50 40 30 open circuit

20 10

open circuit

open circuit

6000 v1 PO43−

pH0 4.0 v2 PO43−

v4 PO4

pH0 8.2 pH0 10.0

3−

Ca4H(PO4)3.3H2O

CaHPO4;

15

Hydroxyapatite

CaHPO4.2H2O

5 0 -5 -10 -15 4

5

6

7

8

9

10

11

12

13

14

pH

(C)

(D)

♣ : Na2SO4 ♥ : HAP

♣ ♥

v3 PO 43−

4000 3000 2000

pH0 4.0 ♣ ♥

♣ ♥

♣

pH0 8.2 pH0 10.0

♣ ♥ ♥

♣ ♣ ♥ ♣

1000 0 400

(B)

10

Intensity (a.u.)

5000

Ca3(PO 4)2 (am 2)

Ca3(PO4)2 (beta);

pH0 10.0

pH0 8.2

pH0 4.0

Ca3(PO4)2 (am 1);

20

-20

0

Intensity (a.u.)

25

Supersaturation index

70

Removal efficiency (%)

(A)

Ca P

♣

ACP ACP 500

600

700

800

900

1000 1100 1200

10

20

30

Raman shift (cm-1)

40

50

60

70

2 theta (degree)

10

(E) 9 pH0 10

8

pH

7 pH0 8.2 6 Thin lines refer to open circuit Thick lines refer to 20 mA

5

pH0 4.0

4 3 0

240

480

720

960

1200

1440

Time (min)

554 555 556 557 558

Figure 1. (A) Effects of initial pH on P and Ca removal efficiency. (B) Supersaturation index calculated from Visual MINTEQ. (C) Raman and (D) XRD patterns of recovered solid products. (E) Change of solution pH in open and closed circuit. Conditions: [Ca(NO3)2·4H2O]=1.0 mM; [Na2SO4] = 50 mM;[Na2HPO4] = 0.6 mM; Current = 20 mA, Time = 24 hours.

25



559

Page 26 of 28

Figure 2. 18.0

11

17.5

(A)

10

16.5 16.0

bulk solution pH measured local pH

15.5

pH

pH

9 8

1 hour 2 hours 4 hours 8 hours 24hours

17.0

(B)

15.0 14.5

7

14.0

6

13.5 13.0

5

12.5

0

560 561 562 563

240

480

720

960

Time (min)

1200

1440

0

100 200 300 400 500 600 700 800 900 1000

Layer tickness (µm)

Figure 2. (A) The measured and (B) theoretically calculated local pH. Conditions: [Ca(NO3)2·4H2O] =1.0 mM; [Na2SO4] = 50 mM; [Na2HPO4] = 0.6 mM; Current = 20 mA, Time = 24 hours.

26


Page 27 of 28


564

Figure 3.

Intensity (a.u.)

(A)

Day 7 Day 4

(C)

Day 1 Day 4 Day 7 Day 11

Intensity (a.u.)

Day 11

v1 PO43−

v1 PO43−

Day 1 Standard HAP 10

20

30

40

50

60

920

70

940

960

980

1000

Raman shift (cm-1)

2 theta (degree) 100

Removal efficiency (%)

(D) 80

60 Ca P

40

20

0 0

1

2

3

4

565 566 567 568 569

5

6

7

8

9

10 11 12 13

Time (day)

Figure 3. (A) XRD patterns, (B) SEM images and (C) Raman spectrum of samples collected under different reaction days. (D) Ca and P concentration change with time elapse. Conditions: [Ca(NO3)2·4H2O]=1.0 mM; [Na2SO4] = 50 mM;[Na2HPO4] = 0.6 mM; Current = 20 mA, pH0 = 8.2; Time = 1 day to 11days.

27



Figure 4.

P Ca Mg

P Ca Mg

20

100

70

90

60 50 40

15

30 10 20 5

10

0

571 572 573 574 575

80

(A)

0

0h

24 h

48 h

(B)

80

Efficiency (P %)

70 60 50 40

Removal effciency (%)

Concentration (mg/L)

570

Page 28 of 28

70 60 50 40 Removal effciency Precipitation effciency

30 20 10 0

9

72 h

10

11

12

pH

Figure 4. (A) Concentration change and removal efficiency of P, Ca and Mg in real wastewater by electrochemical precipitation. (B) Removal efficiency of P by conventional chemical precipitation under different solution pH adjusted by sodium hydroxide.

28


Electrochemical Induced Calcium Phosphate Precipitation: Importance

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