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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
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Electrochemical induced calcium phosphate precipitation: importance of local
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pH
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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
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solutions with pH of 4.0 (56.1%), 8.2 (57.4%) and 10.0 (48.4%) after 24 hours of
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reaction time. While P removal efficiencies are not generally affected by bulk solution 1
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pH, the chemical-physical properties of CaP solids collected on the cathode are still
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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
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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
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depleted in 372 years if current mining rates are maintained.1 In addition, considering
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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
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shortage along with P discharge associated eutrophication, has created increased
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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,
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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
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less economically attractive because of the low concentration of Mg2+ in wastewater.15
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Alternatively, calcium phosphate (CaP), which is the mined component in P rock,
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would be a better solution.16, 17 CaP solids can form without adding Ca2+ since there is
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often already sufficient Ca2+ in bodies of water.18 Therefore, P recovery via CaP
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formation and precipitation is a preferred method, and has received a lot of
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attention.17, 19
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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.
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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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because of the presence of organic acids and inorganic carbonates, significant base
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addition is needed in order to increase the bulk solution pH to a certain level that
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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
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Recently, (bio)electrochemical processes were suggested as next generation
75
technologies for treating (in)organic polluted water
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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
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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
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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
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well-documented.29-31 However, the electrochemical assisted struvite formation, like
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chemical precipitation still relies on dosing of costly Mg2+. Moreover, the importance
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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
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precipitation of phosphate salts but none, to the best of our knowledge, has
25
and recognized as an efficient
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investigated the role of local pH in detail. This is because it is difficult to measure the
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local pH directly, as there still are no reliable pH sensors for detecting the
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electrochemically induced local pH at the electrode surface, though there are some
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special designed lab tools.32, 33 Moreover, the importance of local pH was seemingly
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ignored. Some researchers equate bulk solution pH to local pH and therefore just
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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
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experiment. As an example, Wang et al.29 reported the slight increase of pH near the
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cathode from 7.0 to 7.5 as the cause for pure struvite formation in their
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electrochemical system. However, the local pH can be much higher than can
99
measured.33
be
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To the best of our knowledge the electrochemical induced CaP precipitation on the
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cathode has not been reported, in terms of P removal and recovery and at various bulk
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pH. Although CaP coverage of the cathode might seem unwanted, we see this as an
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opportunity to separate P from waste streams with low P concentrations. Therefore,
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the purpose of this study, was to evaluate the efficiency of a single electrochemical
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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
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evaluating the performance of this system at low, higher and high bulk solution pH
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combined with theoretical calculations. Finally, the efficiency and the cost for
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treatment of real wastewater were addressed to evaluate the potential for this new P
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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
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monohydrogen phosphate (Na2HPO4) and sodium sulphate anhydrous (Na2SO4) were
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purchased
115
(Ca(NO3)2·4H2O) was received from Merck (Germany). Electrodes were provided by
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MAGNETO Special Anodes BV (Schiedam, The Netherlands).
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Electrolysis Setup. The electrochemical cell consisted of two compartments, one
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working cell (500 mL) for CaP precipitation and one tank cell (500 mL) for mixing
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and sampling. The total solution in the two compartments (1000 mL) is circulated
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with a pump at a flow rate of 100 mL/min. The anode material is platinum coated (50
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g/m2) titanium mesh with a round shape (Ø 10 cm, thickness 0.1cm) and it is
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perpendicularly welded to a 10 cm long Ti rod (Ø 0.3 cm). The cathode is a pure
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titanium plate similarly welded (grade A, Ø 8.2 cm, thickness 0.1 cm). A pH sensor
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was placed in the sampling tank to record bulk solution pH change. In some cases, the
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pH electrode (Ø 1.2 cm, Endress Hauser, Germany) was also placed near the cathode
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(~ 1.0 mm), in order to record local pH. The pH sensors were calibrated weekly. The
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diagram of the set sup is shown in Figure S1.
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Electrolysis Experiments. The electrochemical precipitation process was conducted
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with 0.6 mM P and 1.0 mM Ca under constant current (20 mA) conditions and at
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constant ionic strength mediated by 50 mM Na2SO4. The choice for a sulfate salt was
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made because it does not interfere with the precipitation of CaP and does not produce
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harmful chlorine gas as well. While the initial Ca concentration is close to its natural
from
VWR
(Leuven,
Belgium).
Calcium
nitrate
tetrahydrate
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concentration, the initial P concentration was higher compared to real wastewater in
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order to collect sufficient solid samples for further characterization. Where
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appropriate, the bulk solution pH was adjusted by concentrated NaOH or HNO3.
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Unless specified, the electrolysis process was open to air and lasted for 24 hours at
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room temperature. The bulk solution pH was monitored during the whole process and
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logged by a computer program (Liquisys M CPM 253, Endress + Hauser, Naarden,
139
The Netherlands).
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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
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disturb CaP precipitates at the cathode surface, for the sake of solid characterization.
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Then the electrode with precipitates on its surface was air dried. After drying, CaP
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solids were harvested by light scraping. After sampling, the cathode was immersed in
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a 1.0 M HNO3 solution to remove any CaP remaining and then thoroughly rinsed with
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Milli-Q water and dried again for use.
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Analytical methods. The concentrations of P and Ca ions were analyzed by ICP-AES.
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X-ray diffraction (XRD) was used to identify the crystal structure (or absence thereof
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if amorphous) and collected on a Bruker D8 advanced diffractometer equipped with a
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Vantec position sensitive detector and with a Co Kα radiation (λ = 0.179 nm) over a
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range of 5−90° in 0.02 step sizes with an integration time of 0.5 s. Raman spectra
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were obtained using a LabRAM HR Raman spectrometer from Horiba Jobin Yvon to
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obtain bonding information of collected solids. This system is equipped with a
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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
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sample with an Olympus Bx41 microscope equipped with a 50x objective lens, which
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gives a spot size ca.1-2 µm and resolution of 6 cm-1. The detector is a Synapse
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multichannel air cooled (-70°C) CCD. The applied laser power was between 5 and 50
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mW (using density filters). The measurement time varied 5 to 30 s. Finally, the data
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were processed with LabSpec software. The morphology of collected products and
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their elemental compositions were examined by a Scanning Electron Microscope
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(SEM) coupled with Energy dispersive x-ray spectroscopy (EDS) (JEOL-6480LV,
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JEOL Ltd., Japan). Samples were coated with gold using a JEOL JFC-1200 Fine
164
coater at 10 Pa for 30 s.
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Calculations. The degree of saturation (Ω) and saturation index (SI) of a solution
166
regarding a mineral phase, are defined as follows:35
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Ω =
168
SI = log ()
169
Where IAP refers to the ion activity of the associated lattice ions and Ksp is the
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thermodynamic solubility product. The computer program visual MINTEQ36 was
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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
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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
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consumed by other occurring reactions and was homogenously mixed in the local
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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:
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( ) =
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.
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RESULTS AND DISCUSSION
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Effects of initial bulk pH (pH0). As a proof of principle, recovery of P in the
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electrochemical system was evaluated at three pH values including background
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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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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
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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
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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
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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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without rinsing.
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While it is not surprising that P was removed in an alkaline solution, the high removal
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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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precipitation should then be attributed to the production of OH− by electrochemical
244
mediated water reduction at the cathode:
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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.
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xCa2+ + yHPO42−/PO43− + nH2O → ACP↓
(7)
286
Ca2+ + HPO42− + 2H2O → DCPD↓
(8) 13
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8Ca2+ + 4PO43− + 2HPO42− + 5H2O → OCP↓
288
10Ca2+ + 6PO43− + 2OH− → HAP↓
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(9) (10)
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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
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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
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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
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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
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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
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refers to the P removal after filtration through 0.45 µm membrane and therefore this
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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
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co-funded by the Dutch Ministry of Economic Affairs and Ministry of Infrastructure
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and Environment, the European Union Regional Development Fund, the Province of
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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
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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.
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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
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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
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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.
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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
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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.
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