*3.3. Phosphate Recovery*

### 3.3.1. Release of Phosphate from the HTL Solid Phase

Section 3.2 has identified the HTL solid phase as a phosphate-containing source. The release of phosphate from mineral phases by wet chemical extraction (also known as leaching) is the most common technique, as it offers high efficiency and low energy demand. On the other hand, the consumption of the leaching agen<sup>t</sup> and the co-dissolution of inorganic species other than phosphate salts can negatively affect the extraction. The leaching study was carried out with the aim of identifying optimal conditions for the production of phosphate-rich leachate with high purity for struvite precipitation.

**Figure 3.** Different forms of phosphate as well as the crop-available form of phosphate in HTL solid phase (MV ± SD of two replicates; error bars are not shown in the figure). The percentage was calculated as follows: (AP(IP, OP, or CAL-P)/TP) × 100%.

Figure 4a indicates that sulfuric and hydrochloric acids provided a considerable phosphate extraction capacity for the HTL solid phase from both PSS and *SPR* (equilibrium pH of leachate <2). The application of NaOH as a leaching agen<sup>t</sup> was limited. The environmentally beneficial citric acid demonstrated a lower efficiency than that of mineral acids for phosphate recovery from HTL-PSS-SP. The reason for this lower efficiency could be the precipitation of secondary nonapatite phosphate in the case of citric acid. Together with calcium phosphates, other acid-soluble compounds that present in HTL-PSS-SP, such as CaCO3 and Mg-, Fe-, and Al-containing compounds, might be decomposed as well. The released from calcium phosphate PO4<sup>3</sup>− ion can instantly bind with available aluminum or iron ions that have a high affinity for phosphate and precipitates, such as secondary Aland Fe-phosphate salts. According to a study of phosphate solubility by Stumm and Morgan [76], when decreasing the pH, the Ca-phosphate dissolved first, with the Al-phosphate and Fe-phosphate following, respectively. Almost complete acidic phosphate dissolution can be expected at pH < 2 [77,78]. Thus, the precipitation of secondary Al- and Fe-phosphate salts may explain why weaker citric acid (equilibrium pH of leachate >2) exhibited a poorer extraction performance compared to sulfuric and hydrochloric acid for HTL-PSS-SP. This suggestion was also supported by the higher extraction capacities of aluminum and iron ions with H2SO4 than with citric acid. For sulfuric acid as a leaching agent, the extraction capacities of aluminum and iron were 5 mg·g<sup>−</sup><sup>1</sup> and 20 mg·g<sup>−</sup>1, respectively. For citric acid, the capacities were 2 mg·g<sup>−</sup><sup>1</sup> and 10 mg·g<sup>−</sup>1, respectively. In contrast, the extraction efficiency of citric acid for HTL-*SPR*-SP was comparable to those of sulfuric and hydrochloric acids. This finding can be linked to the higher solubility of the phosphate forms in HTL*-SPR*-SP compared to the phosphate forms in HTL-PSS-SP (Figure 3). Furthermore, the contents of aluminum and iron in the HTL solid phase from *SPR* were lower (0.2 wt% and 0.7 wt%, respectively) than from PSS (1.8 wt% and 1.5 wt%, respectively). The molar relation between metal ions and phosphate in the HTL solid phase from *SPR* (Al/P ~0.03 and Fe/P ~0.07) was lower than that from PSS (Al/P ~0.21 and Fe/P ~0.08), which can result in less co-precipitation of secondary Al- and Fe-phosphate salts [77]. While citric acid provided a high extraction capacity for HTL*-SPR*-SP, its application may be limited by its negative effect on subsequent precipitation [79]. Figure 4a conveys that sulfuric acid is more selective to phosphate

release compared to other leaching agents. Significantly less calcium was found in the extract with the application of sulfuric acid, which could relate to the simultaneous co-precipitation of calcium sulfate (gypsum) [80]. This co-precipitation can be seen specifically in the decreasing calcium ion concentrations with increasing concentrations of sulfuric acid in Figure 4b. Large-scale use of sulfuric acid is beneficial from an economic perspective because it offers a low cost as a byproduct of the desulfurization of natural gas and petroleum. Consequently, sulfuric acid was selected as the leaching agen<sup>t</sup> for the following study.

**Figure 4.** Effect of the leaching agen<sup>t</sup> (**a**) and H2SO4 concentration (**b**) on the extraction capacity of elements P, K, Mg, Ca, Fe, and Al for the HTL solid phase from PSS (A) and *SPR* (B).

Phosphate was extracted at multiple sulfuric acid concentrations. Figure 4b indicates that an acid concentration of 0.1M provided incomplete extraction (22% in the case of HTL-PSS-SP and 47% in the case of HTL-*SPR*-SP). The maximum extraction capacity from the HTL solid phase was achieved at 0.5M. Higher acid concentrations, which imply an increase of H<sup>+</sup> per g of the HTL solid phase, did not result in a significant improvement in extraction results. The estimation of acid consumption for the phosphate release was essential for the technical feasibility of phosphate recovery technologies. The amount of acid that is required depends on the chemical composition of the HTL solid phase. Calcium phosphate is its main constituent (Figure 3) and reacts essentially with acid; thus, it is the main acid consumer. The literature has reported an average acid consumption for solid residuals that are rich in calcium phosphate was around 3 mol H<sup>+</sup> pro mol P [77,80,81]. If one assumes an average consumption of sulfuric acid of 3 mol H<sup>+</sup> pro mol P to dissolve the phosphates in the HTL solid phase at a given liquid-to-solid ratio (10:1), then sulfuric acid with concentrations of 0.5M and 0.3M should be sufficient to release phosphate from HTL-PSS-SP and HTL-*SPR*-SP, respectively. These calculated concentrations are positively reflected in the experimental data in Figure 4b.

In the context of sustainability, the potential to recycle the remaining acid-insoluble solid residue should also be considered. This acid-treated residual contains low concentrations of phosphorus. The concentrations of other major elements are also altered. The byproduct of leaching of the HTL solid phase from PSS could be used, for example, for the production of activated carbon [21] or as pozzolan in concrete [82]. Of course, the high sulfur content due to gypsum precipitation in the leaching step (with sulfuric acid) warrants attention. Further work with acid-washed HTL residues is required to improve the current understanding of this material.

### 3.3.2. Phosphate Precipitation in the Form of Struvite

This section examines phosphate separation in the form of struvite (MgNH4PO4·6H2O) by mixing phosphate-rich leachate with the HTL liquid phase that contains an ammonium ion. Ammonium nitrogen (NH4-N) amounted to 320 mg·L−<sup>1</sup> in HTL-PSS-LP and to 6800 mg·L−<sup>1</sup> in HTL-*SPR*-LP. The measured concentration can be lower than the real one because of the possible loss of some ammonium during the thawing of the liquid samples. Since half of the phosphate from *SPR* (Figure 2) remains after HTL in the liquid phase, the possibility of direct struvite crystallization from the HTL liquid phase was examined. Table 2 presents the nutrient distribution in process streams that circulated in nutrient recovery (Figure 1), while Table 3 illustrates the performance of the nutrient recovery.


**Table 2.** Nutrient distribution in the streams circulated by phosphate precipitation.

anot analyzed.

The release rate of phosphate from the HTL solid phase from PSS was lower than that from HTL solid phase from *SPR*. The previous section has suggested that approximately 3 mol H<sup>+</sup> pro mol P was required to dissolve the phosphate from the HTL solid phase from PSS. To set and maintain pH 2, 2.7 mL H2SO4 pro g of the HTL solid phase from PSS was piped into the system. Assuming complete dissociation of H2SO4 corresponded to approximately 1.6 mol H<sup>+</sup> pro mol P at the given liquid-to-solid ratio. There does not seem to be sufficient H<sup>+</sup> to provide complete dissolution of primary and secondary phosphate. In the case of *SPR*, the H2SO4 guided into the system was equivalent to approximately 2 mol H<sup>+</sup> pro mol P. It was sufficient to result in the dissolution of a major part of the phosphate and conforms to the previously specified considerations.


**Table 3.** Performances of phosphate recovery.

The recovery rate of phosphorus from the mix solution was approximately 99% and 66% for PSS and *SPR*, respectively, and approximately 99.9% for direct precipitation from the HTL liquid phase of *SPR.* The pH, molar ratio of the participating ions (PO4<sup>3</sup><sup>−</sup>, Mg<sup>2</sup>+, NH4+), and presence of foreign ions (e.g., Ca2+) are among the major parameters that affect the struvite crystallization [83]. The formation of struvite occurs with the creation of supersaturation (index of the deviation of a dissolved salt from its equilibrium), which is the driving force of crystallization. Supersaturation may be achieved by increasing any or all concentrations of ammonium, magnesium, phosphate, and pH in the solution. In general, pH 9 is optimal for struvite precipitation [84]. To attain oversaturation in the mixed solution (spontaneous formation of struvite), the phosphate-rich leachate was mixed with the HTL liquid phase that was high in ammonium ions in a 1-to-6 volume ratio. The results were the molar ratios of 1.1 and 8.4 for PSS and *SPR*, respectively. The molar ratio of NH4<sup>+</sup>:PO4<sup>3</sup>− in the HTL process water of *SPR* was 13.9. The excess of ammonium is beneficial for struvite crystallization [85,86] and could positively affect the purity of the precipitate, as supported by the fact that the nitrogen content in the precipitate from *SPR* (5.3% and 5.2%) was higher than from PSS (3.9%) and within range of the theoretical value of 5.7%. In contrast, the overdose of ammonium resulted in a low recovery rate of ammonium ions (only 19% and 8% for *SPR* compared to 79% for PSS). The magnesium-to-phosphate molar ratio was also key during struvite crystallization. The magnesium content in leachate and process water was low relative to the phosphate content. Thus, the Mg<sup>2</sup><sup>+</sup>:PO4<sup>3</sup>− ratio had to be adjusted by the addition of a magnesium source, which provided the Mg<sup>2</sup>+ that was required for oversaturation and offsets the negative effect of the calcium ion, which competed with the magnesium ion for the phosphate ion [83]. In the case of PSS, the Mg<sup>2</sup><sup>+</sup>:PO4<sup>3</sup>− ratio was adjusted to approximately 2. The underdose and overdose scenarios were compared for *SPR* as feedstock with a low concentration of Ca2+. In the case of an underdose of magnesium, the recovery rate was approximately 66%, which indicated that the magnesium ion was a limiting factor for struvite precipitation. This finding was in line with reference [86]. Thus, to improve the performance of phosphate recovery for *SPR*, a higher magnesium ion dose was necessary. However, it is notable that the optimal design of phosphate recovery entails a compromise between the performances and chemical consumption.

The percentage of each element (Table 2) in the precipitate was compared with the theoretical value for pure struvite as a reference (12.6 wt% P; 5.7 wt% N; 9.9 wt% Mg). It can be concluded that the precipitate correlated well with struvite. This suggestion was further confirmed by the XRD analysis. The XRD patterns of the precipitate matched the reference struvite (see Figure S3). Figure 5 presents the SEM image of the precipitate that was obtained. The SEM images revealed coarse, irregularly shaped crystals of various sizes. The most commonly observed crystals in the precipitate from PSS had an average length of 10 μm. The crystals were larger in the case of *SPR*. An excess of ammonium

ions may account for the larger size, as already illustrated by other research [83,87]. In addition to the struvite crystals, other solid precipitates were found (marked with yellow arrow). This co-precipitates might be amorphous calcium phosphate. The smaller amount of Ca2+ in the case of *SPR* can result in fewer impurities in the form of calcium phosphate, as evident in the SEM images.

**Figure 5.** Scanning electron microscopy images of the precipitates that were obtained by mixing phosphate-rich leachate from the HTL solid phase with the HTL liquid phase (PSS and *SPR*) and by direct precipitation from the HTL liquid phase (d*SPR*). Solid precipitates different from struvite are marked with a yellow arrow.

It has been indicated that acid dissolving of phosphate followed by precipitation of struvite is an effective approach for HTL-based phosphorus recovery. To scale up and optimize the performance of this approach in terms of the quantity of struvite that it generates, the characteristics (size and purity) of the precipitate, and the consumption of chemicals, it is necessary to gain additional insight into struvite formation, which requires a detailed study.

### 3.3.3. Overall Consideration of Process and Mass Flow of Macronutrients

The phosphate recovery performance in the laboratory-scale study and the elemental balance that was calculated for the HTL pilot plant were used to calculate the mass flow diagram of macronutrients during the HTL coupling with nutrient recovery (see Figure 1). The calculated mass balance can help to reduce the process development time and identify future research potential.

The mass flow for PSS (Figure 6a) implies a relatively low recovery rate of phosphate from unprocessed PSS in the form of struvite. It can firstly be linked to the recovery of a considerable amount of phosphate in the HTL oil phase and, secondly, to the non-optimal efficiency of the leaching step. The increase of acid consumption (Section 3.3.2) to approximately 3 mol H<sup>+</sup> pro mol P might lead to the complete dissolution of the phosphate and an improvement in leaching efficiency. Moreover, the release of more phosphate may result in an increase in the HTL liquid phase and the Mg source consumption to cover the corresponding ion ratio as well as in the NaOH consumption for adjusting the pH. For example, an increase of 1M H2SO4 up to 1 L·h−<sup>1</sup> (corresponds to 3 mol H<sup>+</sup> pro mol P) can require an increased HTL liquid phase to provide NH4<sup>+</sup>:PO4<sup>3</sup>− of 1.6 up to 55 L·h−1. This amount may conflict with the amount of the liquid phase that originates from HTL (calculated at approximately 57 <sup>L</sup>·h−1). The increase in chemical consumption that [88] has been identified as a major part of struvite production costs could result in heightened operating costs. The improvement of nutrient performance necessitates a trade-off between the amount and quality of struvite and the consumption of ammonium and magnesium sources as well as NaOH.

**Figure 6.** Overall process considerations and mass flow of macronutrients for HTL of PSS (**a**) and *SPR* (**b**,**<sup>c</sup>**). The sample ID can be found in Figure 1.

Figure 6b,c presents macronutrients flow from *SPR* in the precipitation struvite from the mix solution of leachate and the HTL liquid phase and in the direct precipitation from the HTL liquid phase. Approximately 40% and 54% of the phosphate from the unprocessed *SPR* was recovered in the struvite in Figure 6b,c, respectively. The HTL liquid phase from *SPR* is the ammonium-rich stream, and an excess of ammonium ions presents in both cases, which can cause a low recovery rate of ammonium ions. In the first case, the amount of ammonium that is required for struvite precipitation can be regulated and optimized by adjusting the mix ratio of leachate and the HTL liquid phase. Meanwhile, in the case of direct precipitation, it is more di fficult to control and adjust the optimal NH4 +:PO4 3−. Further research should consider how to approach the ammonium-rich post-precipitation liquid phase. Possible strategies include stripping ammonia and recovering it in the form of the fertilizer ammonium sulfate [89] or using activated carbon as a sorbent for ammonium separation. In such a case, the struvite crystallization could be beneficial as a pre-treatment technique.

In summary, the recovery of phosphate from the HTL residual stream was successfully performed. The process variations (H2SO4, NaOH, and Mg dosages as well as the amount of the HTL liquid phase) and the composition of the initial solution are the main challenges in developing an e fficient and cost-e ffective process design for the struvite precipitation.
