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New insights into the electrochemical conversion of CO2 to oxalate at stainless steel 304L cathode
Journal of CO₂ Utilization 36 (2020) 105–115
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New insights into the electrochemical conversion of CO2 to oxalate at stainless steel 304L cathode
T
Siddhartha Subramanian, K.R. Athira, M. Anbu Kulandainathan, S. Senthil Kumar*, R.C. Barik* CSIR- Central Electrochemical Research Institute, Karaikudi, Tamilnadu, 630003, India
A R T I C LE I N FO
A B S T R A C T
Keywords: Carbon dioxide Electroreduction Oxalic acid SS 304L alloy Sacrificial anode
Electrochemical conversion of CO2 to value added chemicals is a promising route for its utilization and mitigation from atmosphere. Nevertheless, an industrial process involving CO2 electroreduction is yet to be a reality and requires further studies on improving the existing processes. In the present work, we report new insights into the electrochemical conversion of CO2 to zinc oxalate at SS 304 L cathode in acetonitrile in the presence of a sacrificial Zn anode. The influence of current density, water content in the electrolyte solution and pressure on product selectivity were evaluated. In addition, oxalic acid synthesis from zinc oxalate was studied using ester hydrolysis process. Our results showed that the product yield is affected due to CO2 reduction to carbonate, decomposition of acetonitrile to cyanide and glycolate formed due to partial reduction of oxalate. Experiments performed in a batch reactor at 2 bar pressure showed that an average Faradaic efficiency of 73.9% can be obtained for zinc oxalate. A yield of 58.1% was obtained for the extraction of oxalic acid. Both zinc oxalate and oxalic acid obtained from this process were compared with commercially available products to confirm its purity.
1. Introduction With increasing carbon dioxide (CO2) levels in the atmosphere and growing demands to minimize its release, there has been a remarkable upsurge of interest in improving the existing technologies for its capture and storage [1–3]. There are different approaches for reducing CO2 emissions from the atmosphere. One of the widely studied methods is its conversion to fuels or commodity chemicals using electrochemical reduction at electrodes through heterogeneous catalysis [4–6]. CO2 being a thermodynamically stable molecule, requires a significant amount of energy and electrochemical conversion provides the energy required to accomplish this herculean task. Some of the merits of using electrochemical methods is that it can be performed at room temperature and scaling up the process is not an arduous task [7]. Extensive research work has been done over the years on electrochemical reduction of CO2 to various products on metal electrodes [8]. These studies have elucidated the influence of electrocatalysts and factors such as catholyte composition, current density, pressure and temperature on the selective conversion to different products [9–11]. The fundamental requirement for electrochemical reduction of CO2 to various products in liquid electrolytes is to chose a solvent having a high solubility of CO2. The solubility of CO2 in water at ambient conditions is quite low (0.030 M) [12]. This makes it difficult for bulk scale ⁎
synthesis of products in aqueous systems and increasing the solubility by either increasing the pressure or decreasing the temperature are frequently employed. Many studies have reported increased efficiencies in the electrochemical reduction of CO2 in water at pressures of about 20–30 bar [13–15]. Extensive research work has been done on CO2 reduction in aqueous electrolytes on various metal electrodes for the production of a range of products such as methanol, ethanol, formate and hydrocarbons. The formation of a particular product has been found to be primarily depend on the bonding strength between metal and carbon species, CO2 or the adsorbed CO [16,17]. The other drawback in addition to poor solubility in aqueous electrolytes is the competing hydrogen evolution reaction (HER) at the cathode. This makes it necessary to select suitable catalysts that can suppress HER and favour CO2 reduction. Oxalic acid is widely used as a bleaching agent and in extractive metallurgy [18]. Utilizing CO2 to synthesis oxalate from which oxalic acid can be extracted can be a sustainable process for its conversion into a valuable product. However, the conversion of CO2 to oxalate is not favourable in water due to the presence of H+ ions and is achieved in aprotic solvents at inert electrodes like lead, stainless steel and nickel [19,20]. Acetonitrile, dimethyl formamide, dimethyl sulfoxide has been widely used as solvents for the conversion of CO2 to oxalate with tetra
Corresponding authors. E-mail addresses: [email protected] (S. Senthil Kumar), [email protected] (R.C. Barik).
https://doi.org/10.1016/j.jcou.2019.10.011 Received 13 August 2019; Received in revised form 17 October 2019; Accepted 17 October 2019 2212-9820/ © 2019 Elsevier Ltd. All rights reserved.
Journal of CO₂ Utilization 36 (2020) 105–115
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Fig. 1. Electrochemical setup for the conversion of CO2 to Oxalate in a batch reactor at 2 bar pressure.
2. Experimental
alkyl ammonium salts as the electrolyte. Acetonitrile (ACN) in particular is widely used in electrochemical studies because of its excellent solvation properties for a variety of electrolytes [21]. The advantage of using aprotic solvents is that it eliminates the issue of hydrogen evolution reaction (HER) resulting in higher current efficiencies at atmospheric pressure. Moreover, the solubility of CO2 in these solvents is about 8–9 times higher than in water as it has been reported by Tomita et.al [22] who calculated the solubility in ACN (with traces of water) to be about 0.270 M. This plays a key role in the increased availability of CO2 molecules for electroreduction resulting in higher current efficiencies. In electrochemical reduction of CO2 in aprotic solvents, CO2 is first reduced to highly energetic CO2− radical anion. This radical intermediate ion dimerizes with another CO2− ion by C–C coupling to form oxalate (C2O42−). However, nucleophilic coupling with another CO2 molecule can also occur competitively to produce carbon monoxide (CO) and carbonate (CO32−) which is not of much interest [23]. These two possible reactions that can take place are given in Eqs. (2) and (3). The reaction pathway mostly depends on the catalyst, concentration of dissolved CO2 and the current density [3].
CO2 (l) + e− → CO2− (ads )
(1)
CO2− (ads ) + CO2− (ads ) → C2 O42 −
(2)
CO2− (ads ) + CO2 (l) + e− → CO (g ) + CO32 −
(3)
2.1. Materials Stainless Steel 304 L (SS 304 L) alloy and Zinc plates were purchased from Sigma Aldrich. ACN (HPLC grade, 0.008% water), was purchased from Fischer Scientific. TBAP (99.9%) was purchased from Alfa Aesar and used as obtained. Commercial zinc oxalate was purchased from Otto and is expressed as ZnC2O4-C. Zinc oxalate obtained from electrochemical reduction of CO2 is expressed as ZnC2O4-E. Ethanol and concentrated sulphuric acid (96%) were used for the extraction of oxalic acid from oxalate.
2.2. Electrochemical cell setup Zinc oxalate was synthesized in a single compartment cell with two electrode setup in ambient pressure and compared with the batch reactor at 2 bar pressure. The two electrodes were SS 304 L as cathode and Zn as the anode. The 100 mL cell used was airtight with a gas inlet to pass N2 and CO2 through the solution. Zinc anode was cleaned using dil. HCl (1 N), washed and dried before use. ACN (50 mL) with 0.2 M TBAP was used as the electrolyte solution. Once the electrolyte solution was prepared, electrodes were placed into the cell and N2 gas was purged for 30 min to maintain an oxygen free electrolyte. Thereafter, CO2 was purged for 30 min followed by electrolysis at different current densities (10–30 mA/cm2). The geometrical surface area of both electrodes were kept fixed at 10 cm2 in these experiments. A magnetic pellet was used to continuously stir the solution during the experiment to enhance mass transfer of dissolved CO2 to the electrode surface and electrolysis was carried out for 1 h. Similarly, for the electrolysis carried out in the batch reactor at 2 bar pressure zinc and SS 304 L plates of area 120 cm2 (12 cm x 10 cm) were used. Electrolysis was performed for 8 h. 0.8 L of ACN with 0.2 M TBAP were prepared and fed into the reactor. The distance between the two electrodes was kept fixed at 5 cm and an electric agitator was used for electrolyte stirring. All experiments were performed with a continuous purging of CO2 and electrolyte stirring at 100 rpm. The pressure inside the reactor was maintained at 2 bar throughout the reaction, which was monitored by a pressure gauge installed in the reactor. The schematic of the electrochemical reactor used is shown in Fig.1. The solvent and electrolyte used in the experiment were separated and recovered using
The oxalate anion formed can combine in a galvanic system with a metal ion formed at a sacrificial anode to give metal oxalate. The present work focuses on the previous study of Fischer et al. [24] on electrochemical synthesis of zinc oxalate from CO2 in a one compartment electrochemical cell in ACN solution with 0.2 M tetrabutyl ammonium perchlorate (TBAP) as the electrolyte. Zinc was used as sacrificial anode and stainless steel 304 L (SS 304 L) alloy was used as cathode. The major product obtained which is zinc oxalate is in good agreement, however new observations were found in our study that provides key insights into the formation of carbonate, cyanide and glycolate during the electroreduction process which affects the synthesis of oxalic acid from zinc oxalate. The influence of current density, water content in the electrolyte and pressure on product formation have been investigated and reported here. In addition, the detailed process used for the extraction of oxalic acid from zinc oxalate is discussed. 106
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a rotary evaporator. The solid product obtained in the reaction was filtered, washed with ACN and dried. This product along with concentrated sulphuric acid (96%) and ethanol were used for the extraction of oxalic acid. 2.3. Product characterization Powder X- ray diffractometric analysis (XRD) was performed on a Bruker diffractometer with a Cu anode (Cu Kα, λ = 1.54 Å) using θ- 2θ geometry. Values of 2θ ranged from 100 to 800 and the scan rate employed was 60 per min. Fourier transform infrared spectroscopy data was obtained between 4000 and 400 cm−1 using a Bruker infrared spectrometer. 13C NMR measurements were recorder on a Bruker 400 Avance spectrometer. Scanning electron microscopy was used to study the morphology of the catalyst used. Morphological studies of the product were performed using Field emission scanning Electron Microscopy (FESEM) and Transmission electron microscopy (Tecnai™ G2 TF20) at an accelerating voltage of 200 kV. Atomic force microscopy (Nanonics Imaging Ltd) was used for measuring thickness of synthesized nanosheets.
Fig. 2. Comparison of FTIR spectra (a) ZnC2O4-C, and ZnC2O4-E obtained at (b) 5 mA/cm2, (c) 10 mA/cm2 and (d) 15 mA/cm2 in ambient pressure.
3. Results and discussion 3.1. Role of current density and water content in electrolyte on product formation It is well known that electrochemical reduction of CO2 in acetonitrile produces oxalate and carbonate and the product selectivity depends on the cathode material and applied current density [25]. Table 1 shows the experimental results of electroreduction of CO2 to zinc oxalate (ZnC2O4- E) yield at different current densities (5–30 mA/cm2). Fourier transform Infrared (FTIR) spectroscopy was used for identifying the functional groups present in the product and peaks were compared with commercially available zinc oxalate (ZnC2O4- C). The zinc oxalate was formed for all the applied current densities (Fig. 2 and Table 1) and additional products such as glycolate, cyanide and carbonates were formed at 30 mA/cm2 (Fig. 3 and Table 1). The peaks for ZnC2O4-C (Fig. 2a) matches with the characteristics peaks of ZnC2O4-E at 5 mA/cm2 (Fig. 2a), 10 mA/cm2 (Fig. 2b) and 15 mA/cm2 (Fig. 2c). The peak at 1645 cm−1 is due to the presence C] O group of oxalate. In addition, peaks at 1472 cm−1 and 1028 cm−1 were observed which can be attributed to the presence of CO32− and C–O stretches respectively. This confirms that the obtained product contained some amount of zinc carbonate in addition to zinc oxalate. The amount of zinc oxalate present in the crude product was estimated by KMnO4 titration (See Supplementary data) and the Faradaic efficiency of oxalate at different current densities was calculated (See Table 1). The weight of the product shows rise in the yield as current density increases. It can be also seen from Table 1 that as current densities increased from 5 to 30 mA/cm2, the Faradaic efficiency of oxalate increases. This is because as the current density increases, the concentration of the intermediate CO2− radical anion increases leading to an increased recombination rate that forms oxalate. The FTIR peaks for the applied
Fig. 3. FTIR spectra of product obtained at a current density of 30 mA/cm2 in ambient conditions.
current density of 30 mA/cm2 is shown separately in Fig. 3 to show clearly the formation of additional products along with zinc oxalate and carbonate. At a current density of 30 mA/cm2, zinc oxalate formed with a Faradaic efficiency of 81.1%, however the FTIR spectrum of product obtained showed peaks at 1767 cm−1 and 2016 cm−1 in addition to the peaks at 1579.41 cm−1, 1428 cm−1 and 1028 cm−1. The peak at 1767 cm−1 can be attributed to the presence of glycolate which we believe to have been produced due to the partial reduction of oxalate formed in the reaction. The sharp peak observed here at 1579.4 cm−1 can be attributed to the presence of carbonyl (CO2−) asymmetric stretch [26]. In addition to these peaks, a small peak at 2016 cm−1 was observed and this is possibly due to the presence of nitrile group (C^N). This confirms the formation of zinc cyanide at this current density and is in accordance with a previous article by Weixin et al. who had reported the formation of zinc cyanide at potentials above -3.0 V vs Ag/AgCl due to ACN decomposition at stainless steel cathode [27,28]. ACN decomposition generates cyanide ions which reacts with Zn2+ ions coming from the anode to form zinc cyanide. This shows that at 30 mA/cm2, by products such as zinc cyanide and glycolate are formed in addition to zinc oxalate which are unwanted for oxalic acid synthesis. Table 2 summarizes our assignment of the FTIR bands observed during CO2 reduction at SS 304 L in electrolyte solution based on the spectra in Fig.2 and Fig.3. The other insight observed is that increasing the current density does not inhibit the formation of zinc carbonate which can be attributed
Table 1 Experimental results obtained from electrochemical reduction of CO2 at SS 304 L in ACN-0.2 M TBAP in ambient pressure. Exp No.
Current density (mA/cm2)
Cell Potential (V)
Duration of electrolysis (h)
Weight of product (g)
Faradaic efficiency of oxalate (%)
1 2 3 4
5 10 15 30
2.5 3.0 3.5 4.5
1 1 1 1
0.17 0.32 0.46 0.87
42.61 45.96 51.07 81.1 %
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Table 2 Assignment of FTIR bands observed for solid product obtained from CO2 reduction. Wavenumber (cm−1)
Assignment
Compound
1089, 1320 1363, 1472, 1645, 1767 2015 2964
C-O stretch COO− stretch C-C stretch CO32− C=O COO− stretch CN stretch C-H stretch
Zinc carbonate Zinc oxalate Zinc oxalate Zinc carbonate Zinc oxalate Zinc glycolate Zinc cyanide Zinc oxalate dihydrate (adsorbed water molecules) Zinc oxalate dihydrate
1097 1388 1475 1579
3392, 3408
O-H stretch
to the intermediate CO2− anion reacting with CO2 forming CO and CO32− as discussed earlier, See Eqs. (2) and (3) [3]. However, recent studies have found that the formation of carbonate is due to the presence of trace amounts of water in the electrolyte solution. Figueiredo et al. investigated the effect of water content in ACN on CO2 reduction at copper electrodes and reported that the presence of water results in the formation of carbonate and bicarbonate [29]. A more significant finding is that the formation of CO and CO32− happens independently for this system. Hence, the formation of zinc carbonate may be due to trace amounts of water present and the reactions taking place for its formation are shown in Eqs. (4) to (8).
Zn →
Zn2 +
2H2 O +
2e−
+
Based on these results, an optimum current density of 15 mA/cm2 was chosen as there was no formation of byproducts such as cyanides and glycolate for the large scale production of oxalate in a batch reactor operated at 2 bar pressure. 3.2. Electroreduction of CO2 to oxalate in the batch reactor
(4)
2e−
→ H2 +
Fig. 4. Variation of cell voltage during the electrochemical reduction of CO2 to oxalate in the batch reactor at a constant current density of 15 mA/cm2.
2OH−
The schematic of the batch reactor is shown in Fig. 1. It contains provisions for inserting two electrodes separated by a distance of 5 cm and an agitator placed between them. After setting up the electrodes and feeding the electrolyte solution into the reactor, care was taken to ensure the reactor was airtight and CO2 was fed into the reactor for 30 min with continuous stirring for attaining saturation. A constant pressure of 2 bar was maintained inside the reactor. After 30 min, electrolysis was performed by maintaining a constant current density of 15 mA/cm2 and continued for upto 8 h. As the product formation started and continued for longer hours, its accumulation near the electrodes resulted in a reduced area available for CO2 reduction and an increase in the cell voltage. Fig. 4 shows the increase in cell voltage with time when a constant current is maintained. This can be explained due to the fact that the product accumulation near the vicinity of the cathode shields the active sites for the CO2 molecules to be reduced. Further, the accumulated product reduces the migration of electrons coming from anode leading to an increase in resistance of the electrolyte solution. In order to obtain deeper insight into the increase in cell voltage during electrolysis, electrochemical impedance spectroscopy (EIS) was performed in a single compartment cell using SS 304 L as working electrode (area of 2 cm2), Zn as counter electrode and Ag wire as reference electrode. The volume of electrolyte used was 50 mL. Electrolysis was performed at a constant current density of 15 mA/cm2 for 1 h with constant bubbling of CO2. After 1 h, electrolysis was stopped and EIS was performed with respect to open circuit potential (OCP) with frequency ranging from 1 MHz to 100 mHz. This process was repeated until 6 h of electrolysis. An equivalent circuit was fitted using Randles circuit consisting of an electrolyte solution resistance (Rs) in series with a parallel combination of charge transfer resistance (Rct) and a double layer capacitance (Cdl). The obtained values are reported in Table 4. Nyquist plots obtained after every 1 h of electrolysis is shown in Fig.5 (a). The comparison of EIS spectra clearly show an increase in both Rs and Rct with time. The obtained Rs and Rct values were plotted against time of electrolysis and is shown in Fig.5(b). It can be seen from Table 4 that both Rs and Rct values before electrolysis (t = 0) is higher than those obtained after 1 h of electrolysis. This is because once electrolysis starts, CO2 reduction starts facilitating movement of C2O42− and Zn2+
(5)
OH− + CO2 → HCO3−
(6)
HCO3− + OH− → CO32 − + H2 O
(7)
CO32 − + Zn2 + → ZnCO3
(8)
It is therefore crucial to reduce the water content in the electrolyte to prevent the formation of carbonate and obtain oxalate. The removal of trace amounts of water is done using pre electrolysis using two platinum electrodes and applying a constant potential of -1.8 V vs Ag/ AgCl. This removes water present by liberating it as H2 and O2 gas. However, this is a laborious task for the production of oxalate on a large scale as we are dealing with a higher volume of electrolyte solution, typically about 1 L. On the contrary, the use of molecular sieves to remove moisture present in the electrolyte solution can be a better solution and hence, in our studies ACN with 0.2 M TBAP was stored in molecular sieves for 24 h before each experiment. The precise water content in the electrolyte solutions was measured using Karl Fischer titration with an Metrohm 381 K F Coulometer and is reported in Table 3. It can be seen that the amount of water present in 99.9% ACN with 0.2 M TBAP is 623.4 ppm. When this solution is stored in molecular sieves, the water concentration reduces drastically to 37.6 ppm. Although, the formation of small quantities of zinc carbonate could not be avoided, it does not pose a problem during the extraction of oxalic acid from zinc oxalate because zinc carbonate can be removed as salts of concentrated acids which is discussed later in the ester hydrolysis process. Moreover, for longer duration of electrolysis carried in a batch reactor for 8 h, formation of zinc oxalate becomes dominant once all the residual water gets used up in the formation of carbonate. Table 3 Water content of the electrolyte solutions used in the experiments. Solution
Water concentration (ppm)
ACN (99.9 %) + 0.2 M TBAP ACN (99.9 %) + 0.2 M TBAP after removing water using molecular sieves
623.4 37.6
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microscopy and the results obtained are as shown in Fig. 6. It can be seen clearly from these images that there is no significant change in the morphology of the catalyst. The composition characterized by EDX results also reveal no significant change confirming no significant degradation of the elements present in the alloy (Ni and Cr).
Table 4 Solution resistance and charge transfer resistance after various electrolysis duration carried out a constant current density of 15 mA/cm2. Time (hrs)
Solution resistance, Rs (ohm. cm2)
Charge transfer resistance, Rct (ohm. cm2)
Cdl (F/cm2)
0 1 2 3 4 5 6
5.74 4.84 5.20 5.60 6.22 7.13 8.99
4.23 × 104 1.54 × 104 2.28 × 104 2.67 × 104 3.41 × 104 4.00 × 104 5.33 × 104
3.67 × 10−5 3.96 × 10−5 3.29 × 10−5 3.29 × 10−5 2.70 × 10−5 2.31 × 10−5 1.67 × 10−5
3.4. Rate of dissolution of sacrificial anode Another important parameter that has to be considered in this process is the rate of dissolution of the sacrificial zinc anode in the reactor in order to estimate the amount of zinc required for the 8 -h electrolysis. To determine this, the weight of zinc before and after the reaction was measured in each experiment. The initial weight of zinc in the first experiment was 100.13 g and after performing electrolysis for 8 h, the weight of zinc was found to be 82.03 g implying that 18.1 g of zinc was consumed in the reaction. This experimental dissolution rate of zinc was compared with the theoretically predicted dissolution rate of zinc. Pure zinc has a theoretical maximum capacity of 820 mA h/g which is also termed as anode capacity [32]. For 8 h of electrolysis operating at a constant current of 1.8 A (current density of 15 mA/cm2) the dissolution rate was calculated and found to be 17.56 g (See Eq. (11)). This means that theoretically, a minimum of 17.56 g of zinc is required to use it as a sacrificial anode which can continuously supply 1.8 A of current for 8 h. The difference of 0.54 g (18.10 -17.56) obtained in the experiment is due to the zinc operating at around 95% efficiency which means a slightly higher mass of 18.1 g of zinc is required to perform this electrolysis.
ions resulting in a drop in the solution resistance (increase in conductance). However, EIS performed after 1 h of electrolysis show an increase in both Rs and Rct and it increases steadily till 6 h of electrolysis. This is in accordance with the increase in cell voltage obtained in the 8 h of electrolysis performed in the batch reactor at a constant current density of 15 mA/cm2. The reason for this steady rise is essentially due to the accumulation of product (zinc oxalate) formed that blocks the active sites of the catalyst (SS 304 L). Hence, due to an increase in the overall resistance of the system, the cell voltage rises (at constant current density) in accordance with Ohm’s law (V = IR). This phenomenon is similar to potassium carbonate precipitation observed on the cathode side in CO2 electrolyzers using aqueous KOH solution as electrolyte. In such an electrolyzer having an anion exchange membrane, a similar rise in cell voltage is observed due to the precipitated carbonate blocking the active sites of the catalyst [30,31]. It is evident from Fig. 4 that the cell voltage increases steeply after 4 h and reaches a value of 6.7 V at 8 h when product accumulation in the reactor goes very high. This increase in the cell voltage can be overcome by creating a continuous electrochemical process where the electrolyte solution can be allowed to flow through a filtration unit to separate the solid product obtained in the reactor, subsequently recycling the product free filtrate (electrolyte solution) back to the reactor with no/very little solid product in the electrolyte solution. Current efforts are directed towards this task.
Q = I × t = (1.8A × 8 × 3600) = 51, 840 C C = 820mAhg −1 = (820 × 10−3 × 3600) Cg −1 = 2952 Cg −1
w=
51840 C = 17.56g 2952 Cg −1
(9) (10)
(11)
Here, ‘Q’ is the amount of charge applied, ‘C’ is the anode capacity of zinc and ‘w’ is the weight of zinc dissolved. This dissolution rate matched well with the theoretically predicted dissolution rate of zinc. In order to safely operate the electrochemical reduction process, the amount of zinc was kept very high (∼ 100 g) with respect to the minimum 18 g requirement. The total mass of the conversion product obtained from the reactor was found to be 38.1 g. This means that the remaining 20 g (38.1 g18.1 g) comes from oxalate and carbonate ions as the product is a mixture of zinc oxalate and zinc carbonate. This is as expected and importantly no other by-products such as zinc cyanide or glycolate is observed at this current density as discussed earlier. The obtained product was then filtered, washed using ACN, dried and qualitative analysis was done using FTIR spectroscopy which confirmed the presence of zinc oxalate and some amount of zinc carbonate. (See Fig. S2,
3.3. Stability of catalyst The stability of catalyst used is one of the important factors that has to be taken into consideration in electroreduction of CO2. From our studies, the catalyst used (SS 304 L) was relatively stable for 8 h of electrolysis. The increase in cell voltage during electrolysis was due to the accumulation of the solid product in the vicinity of the cathode as explained before and not essentially due to catalyst deactivation. To investigate this, the surface morphology of SS 304 L electrode before and after electrolysis was studied using scanning electron
Fig. 5. (a) EIS spectra of SS 304 L in CO2 saturated ACN-TBAP at open circuit potential after various electrolysis duration. (b) Variation of solution resistance and charge transfer resistance after various electrolysis duration, carried out a constant current density of 15 mA/cm2. 109
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Fig. 6. Scanning electron microscopy images and EDX results of SS 304 L before (a), (b) and after (c), (d) the reaction showing no significant change in the morphology.
Table 5 Experimental results of CO2 electroreduction in the batch reactor with S.S 304 L cathode and Zn anode. Exp No.
Cathode
Current density (mA/cm2)
Cell Voltage (V)
1 2 3 4 5
S.S S.S S.S S.S S.S
15 15 15 15 15
3.5 3.5 3.5 3.5 3.5
304 L 304 L 304 L 304 L 404L
– – – –
6.7 6.6 6.1 7.1
Duration of Electrolysis (h)
Weight of crude product (g)
Faradaic efficiency of oxalate (%)
8 8 8 8 4
38.1 41.2 37.2 35.1 22.4
69.1 66.2 81.1 79.3 77.3
Table 6 Experimental results of the extraction of oxalic acid from zinc oxalate. Exp. No
Weight of zinc oxalate (g)
Volume of ethanol (mL)
Volume of conc. H2SO4 (mL)
Volume of water (mL)
Theoretical weight of oxalic acid (g)
Weight of oxalic acid obtained (g)
Yield (%)
1 2 3 4
2 2 4 4
10 10 20 20
4 4 7 7
15 15 30 30
2.30 2.30 1.15 1.15
0.61 0.65 1.34 1.31
52.9 56.5 58.2 56.9
Supplementary data). The electrolysis experiments for a duration of 8 h were repeated four times and the average weight of solid product obtained was found to be 37.9 ( ± 3.3) g with a Faradaic efficiency of 73.9% ( ± 7.3) for oxalate. The results obtained are shown in Table 5. This increase in the Faradaic efficiency from 51.1% at ambient pressure to 73.9% at 2 bar can be attributed to the increased solubility of carbon dioxide at higher pressure and reduced water content in the electrolyte. The product was then used as the starting material for the synthesis of oxalic acid (See Table 6).
All the marked diffraction peak positions are in good agreement with the standard JCPDS card (Joint committee on powder Diffraction standard) No. 25-1029 confirming that the product is zinc oxalate dihydrate. Both diffractograms show well defined singlet at 2θ = 18.70 for (202) plane of zinc oxalate. X- ray diffraction peak positions of zinc oxalate reported by Guo et.al for the synthesis of rod like zinc oxalate also confirm the same [33].
3.5. X- ray diffraction studies of the obtained product
Fig. 8 shows the XPS of ZnC2O4-C and ZnC2O4-E with high resolution scans for Zn 2p, C 1s, O 1s, peaks. The B.E. value at 1045.1 eV and 1046.1 eV for of Zn 2p1/2 orbital and Zn 2p3/2 orbital is at 1022.5 eV and 1023.5 eV for ZnC2O4-C and ZnC2O4-E respectively [34]. The small shift in B.E values of Zn may be due to difference in the morphology of zinc oxalate as confirmed by FESEM images of both ZnC2O4-C and ZnC2O4-E. Moreover, extra peaks at 1020.8 eV and 1044.1 eV are observed for ZnC2O4-E. This could be attributed to the formation of traces of zinc carbonate in addition to zinc oxalate which is also confirmed by
3.6. XPS analysis of zinc oxalate
In order to confirm the product obtained from electroreduction of CO2 in the batch reactor to be zinc oxalate, X- ray diffraction analysis was performed in addition to FTIR spectroscopy and was compared with ZnC2O4-C. The diffraction pattern obtained is shown in Fig. 7. The XRD patterns acquired clearly show the crystalline structure of zinc oxalate confirming the product obtained in addition to FTIR spectroscopy results. 110
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Fig. 7. X-ray diffraction patterns of (a) ZnC2O4-E obtained at 15 mA/cm2 and (b) ZnC2O4-C.
observed from Fig. 9(a) that ZnC2O4-C composed of micro sheets of different sizes. ZnC2O4-E on the other hand composed of an aggregate of particles, which is evident from Fig. 9(c) and (d). It is also clearly evident from these images that the average particle size of ZnC2O4-C is higher than that of ZnC2O4-E. In order to obtain a clear morphology of ZnC2O4-E and determine its particle size, transmission electron microscopy (TEM) was performed and results obtained are shown in Fig.10. It can be seen that nanosheets of zinc oxalate are formed, having an average length and breadth of 85.2 nm and 32.4 nm respectively. In contrast, the average length and breadth of microsheets of ZnC2O4-C was 461.7 nm and 261.8 nm respectively. This shows that zinc oxalate synthesized from electrochemical reduction of CO2 produces nanoparticles which may be of interest in other applications such as in the synthesis of zinc oxide [37,38]. Further, the thickness of zinc oxalate
FTIR spectroscopy as discussed earlier. Similarly, the high resolution C 1s peaks are qualitatively similar for both ZnC2O4-C and ZnC2O4-C. C 1s peaks for ZnC2O4-C are observed at 288.62 eV and 284.98 eV which can be attributed to the presence of C = O and C-C respectively both of which are present in zinc oxalate and is in close agreement with previously reported results [35,36]. Further, the high resolution O 1s peak is observed at 531.81 eV for ZnC2O4-C and at 531.65 eV for ZnC2O4-E. This small difference in the B.E is possibly due to the traces of zinc carbonate present in addition to zinc oxalate in ZnC2O4-E.
3.7. Morphological studies of zinc oxalate The morphological study of ZnC2O4-C and ZnC2O4-E was carried out with field emission scanning electron microscopy (FESEM). It can be
Fig. 8. High resolution XPS spectra of (i) Zn 2p, (ii) C 1s and (iii) O 1s in (a) ZnC2O4-C and (b) ZnC2O4-E. 111
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Fig. 9. FESEM images of (a) and (b) ZnC2O4-C, (c) and (d) ZnC2O4-E showing a clear difference in their morphology.
Diethyl ether for instance is immiscible in water and forms an organic phase but is not an effective solvent for this extraction process due to the poor solubility of oxalic acid in ethers [39]. In our experiments, calculated amounts of ZnC2O4-E and an excess of 4 N sulphuric acid were taken and mixed vigorously till it completely dissolved. To this solution, calculated amounts of diethyl ether was added and agitated vigorously in a separating funnel which formed two immiscible phases. The organic phase was separated and overnight recrystallization produced crystals of oxalic acid. However, the yield obtained was very low. 2 g of zinc oxalate yielded 35 mg of oxalic acid, which is only about 3.5%. This can be attributed due to the poor solubility of oxalic acid in ethers. In organic solvents such as ethanol and methanol, oxalic acid has a relatively higher solubility [40]. However, these solvents are miscible in water and form azeotropes which makes it difficult to separate it from zinc oxalate.
nanosheets synthesized from CO2 have been probed by atomic force microscopy measurements on mica substrate. The average thickness (Z) was found to be 27.5 nm as depicted in Fig. 11. The average size (X) is in the order of 90 nm (1.1 μm − 0.2 μm) and is in close agreement with the length of nanosheet measured from TEM which is 85.2 nm as mentioned earlier.
3.8. Oxalic acid extraction from zinc oxalate Zinc oxalate obtained from electrolysis experiments was washed, dried and used as the starting material for the extraction of oxalic acid. Zinc oxalate is insoluble in many solvents and dissolves only in strong acids like hydrochloric acid and sulphuric acid. When it is treated with 4 N sulphuric acid, it produces zinc sulphate and oxalic acid (See Eq. (12)). However, both zinc sulphate and oxalic acid are soluble in water. Hence, the use of an organic solvent that is immiscible in water and having a high solubility of oxalic acid in it is essential. It is important that the organic solvent satisfies both these conditions for the effective extraction of oxalic acid.
ZnC2 O4 + H2 SO4 → ZnSO4 + H2 C2 O4
(12)
It is therefore important to separate zinc initially from zinc oxalate since both zinc sulphate and oxalic acid are soluble in water. In the second extraction process that we studied, zinc was first removed from
Fig. 10. TEM images ZnC2O4-E at (a) 200 nm and (b) 100 nm showing the nanosheet morphology of zinc oxalate synthesized from CO2. 112
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Fig. 11. The AFM measured thickness of zinc oxalate nanosheet is 27.5 nm. AFM image is shown with a zinc oxalate nanosheet.
Fig. 12. Flowchart of Ester hydrolysis process used for the extraction of oxalic acid from zinc oxalate.
It is important to note here that some amount of zinc carbonate was also present in addition to zinc oxalate obtained from CO2. However, upon treatment with concentrated sulphuric acid, the zinc carbonate went off as zinc sulphate, releasing carbon dioxide and water which is shown in Eq. (13). The limitation in this process is the release of CO2 and hence its concentration in the product has to be minimized.
zinc oxalate as zinc sulphate using concentrated sulphuric acid (96%) without adding any water. Ethanol with a high solubility of oxalic acid was used as the solvent for oxalic acid extraction. Upon separating zinc sulphate, water was added to the filtrate for performing ester hydrolysis. The schematic of the steps used in this extraction process is shown in Fig. 12. Briefly, ZnC2O4-E was first washed, dried and dispersed in ethanol in calculated proportion and stirred vigorously for 1 h. To this mixture, concentrated sulphuric acid was added and stirred for about 3 h. The white slurry obtained was then centrifuged and the precipitate was separated out which was confirmed to be zinc sulphate. Excess of water was then added to the filtrate solution containing diethyl ester and heated at 700 C for 30 min for the completion of ester hydrolysis. The vapours coming from this reaction were condensed to obtain ethanol which can be reused. Thereafter, the solution was cooled and crystallization was performed multiple times. This yielded white crystals which was confirmed to be pure oxalic acid using 13C NMR spectroscopy.
ZnCO3 + H2 SO4 → ZnSO4 + CO2 + H2 O 13
(13)
C NMR spectroscopy was used for the characterization of oxalic acid crystals extracted from zinc oxalate and was compared with commercially available oxalic acid. For NMR studies, 50 mg of oxalic acid was dissolved in 1 mL of deuterated water (D2O) and the obtained results are as shown in Fig. 13. It can be seen that a single peak at 160.41 ppm is obtained for the extracted oxalic acid and at 161.69 ppm for the commercially obtained oxalic acid showing that the extracted oxalic acid is as pure as commercially obtained oxalic acid. A maximum yield of 58.2% was obtained which is significantly higher than the other process used for separation using diethyl ether. We are currently 113
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Fig. 13. 13C NMR spectroscopy results of oxalic acid, (a) commercially obtained and. (b) extracted from zinc oxalate.
working on methods to improve the yield.
loss encountered in this process.
3.9. Electrolyte and solvent recovery after electrolysis
Declaration of Competing Interest There are no potential conflicts of interest reported by the authors.
One of the main limitations of using non-aqueous solvents for CO2 reduction is its high cost. It is therefore important to recover both the solvent and electrolyte for reuse after the reaction. To study this, the liquid electrolyte used in the reaction was first filtered to separate the solid product. After separation, this liquid was taken in a round bottom flask and loaded into a rotary evaporator. Upon distillation of the solvent ACN, white solid crystals precipitated which was separated and dried. This white crystal was confirmed to be tetrabutyl ammonium perchlorate (TBAP) using FTIR spectroscopy (See Fig.S3, Supplementary data) and a yield of 86.7% was obtained. Both ACN and TBAP can be recovered which becomes an important factor when the economics of this process is considered.
Acknowledgement This work was supported by CSIR Mission mode project on “Catalysis for Sustainable Development (CSD)”, Project No: HCP-0009. The authors are grateful to Central Instrumentation Facility, CSIRCECRI for providing access to FTIR, XRD, SEM and TEM facilities. The authors also thank Mr. Amuthan Dekshinamoorthy for helping us in AFM measurements and Mr. Vitalis Chukwuike for valuble discussions on electrochemical impedance spectroscopy. Appendix A. Supplementary data
4. Conclusions
Supplementary material related to this article can be found, in the online version, at https://doi.org/10.1016/j.jcou.2019.10.011.
In summary, the electrochemical reduction of CO2 to oxalate at SS 304 L cathode in ACN is an efficient method for the production of oxalic acid. Minimizing the water content in the electrolyte solution, operating at an optimum current density of 15 mA/cm2 and maintaining a pressure of 2 bar are the three important factors for obtaining a high yield of zinc oxalate. As the current density is increased to 30 mA/cm2, byproducts like zinc cyanide and glycolate starts to form due to ACN decomposition and partial reduction of oxalate respectively. Although, the formation of small amounts of zinc carbonate cannot be inhibited completely, it does not pose an issue in the oxalic acid extraction process as it can be removed as the salt of concentrated acid. The limitations of this process are the relatively lower yield of oxalic acid (58.1%), release of traces of CO2 during oxalic acid extraction that occurs due to conversion of zinc carbonate to zinc sulphate and the accumulation of product in the reactor that causes an increase in the cell voltage. Average Faradaic efficiency of 73.9% can be obtained for the conversion of CO2 to oxalate and there is a significant increase in the efficiency upon increasing the pressure. ACN and TBAP used in the reaction can be recovered and a recovery rate of 86.1% of TBAP with can be achieved without any byproducts. This becomes a significant factor in process economics due to its high cost. Current efforts are directed towards increasing the yield of oxalic acid extracted from zinc oxalate and designing a continuous process for minimizing the ohmic
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New insights into the electrochemical conversion of CO2 to oxalate at stainless steel 304L cathode
T
Siddhartha Subramanian, K.R. Athira, M. Anbu Kulandainathan, S. Senthil Kumar*, R.C. Barik* CSIR- Central Electrochemical Research Institute, Karaikudi, Tamilnadu, 630003, India
A R T I C LE I N FO
A B S T R A C T
Keywords: Carbon dioxide Electroreduction Oxalic acid SS 304L alloy Sacrificial anode
Electrochemical conversion of CO2 to value added chemicals is a promising route for its utilization and mitigation from atmosphere. Nevertheless, an industrial process involving CO2 electroreduction is yet to be a reality and requires further studies on improving the existing processes. In the present work, we report new insights into the electrochemical conversion of CO2 to zinc oxalate at SS 304 L cathode in acetonitrile in the presence of a sacrificial Zn anode. The influence of current density, water content in the electrolyte solution and pressure on product selectivity were evaluated. In addition, oxalic acid synthesis from zinc oxalate was studied using ester hydrolysis process. Our results showed that the product yield is affected due to CO2 reduction to carbonate, decomposition of acetonitrile to cyanide and glycolate formed due to partial reduction of oxalate. Experiments performed in a batch reactor at 2 bar pressure showed that an average Faradaic efficiency of 73.9% can be obtained for zinc oxalate. A yield of 58.1% was obtained for the extraction of oxalic acid. Both zinc oxalate and oxalic acid obtained from this process were compared with commercially available products to confirm its purity.
1. Introduction With increasing carbon dioxide (CO2) levels in the atmosphere and growing demands to minimize its release, there has been a remarkable upsurge of interest in improving the existing technologies for its capture and storage [1–3]. There are different approaches for reducing CO2 emissions from the atmosphere. One of the widely studied methods is its conversion to fuels or commodity chemicals using electrochemical reduction at electrodes through heterogeneous catalysis [4–6]. CO2 being a thermodynamically stable molecule, requires a significant amount of energy and electrochemical conversion provides the energy required to accomplish this herculean task. Some of the merits of using electrochemical methods is that it can be performed at room temperature and scaling up the process is not an arduous task [7]. Extensive research work has been done over the years on electrochemical reduction of CO2 to various products on metal electrodes [8]. These studies have elucidated the influence of electrocatalysts and factors such as catholyte composition, current density, pressure and temperature on the selective conversion to different products [9–11]. The fundamental requirement for electrochemical reduction of CO2 to various products in liquid electrolytes is to chose a solvent having a high solubility of CO2. The solubility of CO2 in water at ambient conditions is quite low (0.030 M) [12]. This makes it difficult for bulk scale ⁎
synthesis of products in aqueous systems and increasing the solubility by either increasing the pressure or decreasing the temperature are frequently employed. Many studies have reported increased efficiencies in the electrochemical reduction of CO2 in water at pressures of about 20–30 bar [13–15]. Extensive research work has been done on CO2 reduction in aqueous electrolytes on various metal electrodes for the production of a range of products such as methanol, ethanol, formate and hydrocarbons. The formation of a particular product has been found to be primarily depend on the bonding strength between metal and carbon species, CO2 or the adsorbed CO [16,17]. The other drawback in addition to poor solubility in aqueous electrolytes is the competing hydrogen evolution reaction (HER) at the cathode. This makes it necessary to select suitable catalysts that can suppress HER and favour CO2 reduction. Oxalic acid is widely used as a bleaching agent and in extractive metallurgy [18]. Utilizing CO2 to synthesis oxalate from which oxalic acid can be extracted can be a sustainable process for its conversion into a valuable product. However, the conversion of CO2 to oxalate is not favourable in water due to the presence of H+ ions and is achieved in aprotic solvents at inert electrodes like lead, stainless steel and nickel [19,20]. Acetonitrile, dimethyl formamide, dimethyl sulfoxide has been widely used as solvents for the conversion of CO2 to oxalate with tetra
Corresponding authors. E-mail addresses: [email protected] (S. Senthil Kumar), [email protected] (R.C. Barik).
https://doi.org/10.1016/j.jcou.2019.10.011 Received 13 August 2019; Received in revised form 17 October 2019; Accepted 17 October 2019 2212-9820/ © 2019 Elsevier Ltd. All rights reserved.
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Fig. 1. Electrochemical setup for the conversion of CO2 to Oxalate in a batch reactor at 2 bar pressure.
2. Experimental
alkyl ammonium salts as the electrolyte. Acetonitrile (ACN) in particular is widely used in electrochemical studies because of its excellent solvation properties for a variety of electrolytes [21]. The advantage of using aprotic solvents is that it eliminates the issue of hydrogen evolution reaction (HER) resulting in higher current efficiencies at atmospheric pressure. Moreover, the solubility of CO2 in these solvents is about 8–9 times higher than in water as it has been reported by Tomita et.al [22] who calculated the solubility in ACN (with traces of water) to be about 0.270 M. This plays a key role in the increased availability of CO2 molecules for electroreduction resulting in higher current efficiencies. In electrochemical reduction of CO2 in aprotic solvents, CO2 is first reduced to highly energetic CO2− radical anion. This radical intermediate ion dimerizes with another CO2− ion by C–C coupling to form oxalate (C2O42−). However, nucleophilic coupling with another CO2 molecule can also occur competitively to produce carbon monoxide (CO) and carbonate (CO32−) which is not of much interest [23]. These two possible reactions that can take place are given in Eqs. (2) and (3). The reaction pathway mostly depends on the catalyst, concentration of dissolved CO2 and the current density [3].
CO2 (l) + e− → CO2− (ads )
(1)
CO2− (ads ) + CO2− (ads ) → C2 O42 −
(2)
CO2− (ads ) + CO2 (l) + e− → CO (g ) + CO32 −
(3)
2.1. Materials Stainless Steel 304 L (SS 304 L) alloy and Zinc plates were purchased from Sigma Aldrich. ACN (HPLC grade, 0.008% water), was purchased from Fischer Scientific. TBAP (99.9%) was purchased from Alfa Aesar and used as obtained. Commercial zinc oxalate was purchased from Otto and is expressed as ZnC2O4-C. Zinc oxalate obtained from electrochemical reduction of CO2 is expressed as ZnC2O4-E. Ethanol and concentrated sulphuric acid (96%) were used for the extraction of oxalic acid from oxalate.
2.2. Electrochemical cell setup Zinc oxalate was synthesized in a single compartment cell with two electrode setup in ambient pressure and compared with the batch reactor at 2 bar pressure. The two electrodes were SS 304 L as cathode and Zn as the anode. The 100 mL cell used was airtight with a gas inlet to pass N2 and CO2 through the solution. Zinc anode was cleaned using dil. HCl (1 N), washed and dried before use. ACN (50 mL) with 0.2 M TBAP was used as the electrolyte solution. Once the electrolyte solution was prepared, electrodes were placed into the cell and N2 gas was purged for 30 min to maintain an oxygen free electrolyte. Thereafter, CO2 was purged for 30 min followed by electrolysis at different current densities (10–30 mA/cm2). The geometrical surface area of both electrodes were kept fixed at 10 cm2 in these experiments. A magnetic pellet was used to continuously stir the solution during the experiment to enhance mass transfer of dissolved CO2 to the electrode surface and electrolysis was carried out for 1 h. Similarly, for the electrolysis carried out in the batch reactor at 2 bar pressure zinc and SS 304 L plates of area 120 cm2 (12 cm x 10 cm) were used. Electrolysis was performed for 8 h. 0.8 L of ACN with 0.2 M TBAP were prepared and fed into the reactor. The distance between the two electrodes was kept fixed at 5 cm and an electric agitator was used for electrolyte stirring. All experiments were performed with a continuous purging of CO2 and electrolyte stirring at 100 rpm. The pressure inside the reactor was maintained at 2 bar throughout the reaction, which was monitored by a pressure gauge installed in the reactor. The schematic of the electrochemical reactor used is shown in Fig.1. The solvent and electrolyte used in the experiment were separated and recovered using
The oxalate anion formed can combine in a galvanic system with a metal ion formed at a sacrificial anode to give metal oxalate. The present work focuses on the previous study of Fischer et al. [24] on electrochemical synthesis of zinc oxalate from CO2 in a one compartment electrochemical cell in ACN solution with 0.2 M tetrabutyl ammonium perchlorate (TBAP) as the electrolyte. Zinc was used as sacrificial anode and stainless steel 304 L (SS 304 L) alloy was used as cathode. The major product obtained which is zinc oxalate is in good agreement, however new observations were found in our study that provides key insights into the formation of carbonate, cyanide and glycolate during the electroreduction process which affects the synthesis of oxalic acid from zinc oxalate. The influence of current density, water content in the electrolyte and pressure on product formation have been investigated and reported here. In addition, the detailed process used for the extraction of oxalic acid from zinc oxalate is discussed. 106
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a rotary evaporator. The solid product obtained in the reaction was filtered, washed with ACN and dried. This product along with concentrated sulphuric acid (96%) and ethanol were used for the extraction of oxalic acid. 2.3. Product characterization Powder X- ray diffractometric analysis (XRD) was performed on a Bruker diffractometer with a Cu anode (Cu Kα, λ = 1.54 Å) using θ- 2θ geometry. Values of 2θ ranged from 100 to 800 and the scan rate employed was 60 per min. Fourier transform infrared spectroscopy data was obtained between 4000 and 400 cm−1 using a Bruker infrared spectrometer. 13C NMR measurements were recorder on a Bruker 400 Avance spectrometer. Scanning electron microscopy was used to study the morphology of the catalyst used. Morphological studies of the product were performed using Field emission scanning Electron Microscopy (FESEM) and Transmission electron microscopy (Tecnai™ G2 TF20) at an accelerating voltage of 200 kV. Atomic force microscopy (Nanonics Imaging Ltd) was used for measuring thickness of synthesized nanosheets.
Fig. 2. Comparison of FTIR spectra (a) ZnC2O4-C, and ZnC2O4-E obtained at (b) 5 mA/cm2, (c) 10 mA/cm2 and (d) 15 mA/cm2 in ambient pressure.
3. Results and discussion 3.1. Role of current density and water content in electrolyte on product formation It is well known that electrochemical reduction of CO2 in acetonitrile produces oxalate and carbonate and the product selectivity depends on the cathode material and applied current density [25]. Table 1 shows the experimental results of electroreduction of CO2 to zinc oxalate (ZnC2O4- E) yield at different current densities (5–30 mA/cm2). Fourier transform Infrared (FTIR) spectroscopy was used for identifying the functional groups present in the product and peaks were compared with commercially available zinc oxalate (ZnC2O4- C). The zinc oxalate was formed for all the applied current densities (Fig. 2 and Table 1) and additional products such as glycolate, cyanide and carbonates were formed at 30 mA/cm2 (Fig. 3 and Table 1). The peaks for ZnC2O4-C (Fig. 2a) matches with the characteristics peaks of ZnC2O4-E at 5 mA/cm2 (Fig. 2a), 10 mA/cm2 (Fig. 2b) and 15 mA/cm2 (Fig. 2c). The peak at 1645 cm−1 is due to the presence C] O group of oxalate. In addition, peaks at 1472 cm−1 and 1028 cm−1 were observed which can be attributed to the presence of CO32− and C–O stretches respectively. This confirms that the obtained product contained some amount of zinc carbonate in addition to zinc oxalate. The amount of zinc oxalate present in the crude product was estimated by KMnO4 titration (See Supplementary data) and the Faradaic efficiency of oxalate at different current densities was calculated (See Table 1). The weight of the product shows rise in the yield as current density increases. It can be also seen from Table 1 that as current densities increased from 5 to 30 mA/cm2, the Faradaic efficiency of oxalate increases. This is because as the current density increases, the concentration of the intermediate CO2− radical anion increases leading to an increased recombination rate that forms oxalate. The FTIR peaks for the applied
Fig. 3. FTIR spectra of product obtained at a current density of 30 mA/cm2 in ambient conditions.
current density of 30 mA/cm2 is shown separately in Fig. 3 to show clearly the formation of additional products along with zinc oxalate and carbonate. At a current density of 30 mA/cm2, zinc oxalate formed with a Faradaic efficiency of 81.1%, however the FTIR spectrum of product obtained showed peaks at 1767 cm−1 and 2016 cm−1 in addition to the peaks at 1579.41 cm−1, 1428 cm−1 and 1028 cm−1. The peak at 1767 cm−1 can be attributed to the presence of glycolate which we believe to have been produced due to the partial reduction of oxalate formed in the reaction. The sharp peak observed here at 1579.4 cm−1 can be attributed to the presence of carbonyl (CO2−) asymmetric stretch [26]. In addition to these peaks, a small peak at 2016 cm−1 was observed and this is possibly due to the presence of nitrile group (C^N). This confirms the formation of zinc cyanide at this current density and is in accordance with a previous article by Weixin et al. who had reported the formation of zinc cyanide at potentials above -3.0 V vs Ag/AgCl due to ACN decomposition at stainless steel cathode [27,28]. ACN decomposition generates cyanide ions which reacts with Zn2+ ions coming from the anode to form zinc cyanide. This shows that at 30 mA/cm2, by products such as zinc cyanide and glycolate are formed in addition to zinc oxalate which are unwanted for oxalic acid synthesis. Table 2 summarizes our assignment of the FTIR bands observed during CO2 reduction at SS 304 L in electrolyte solution based on the spectra in Fig.2 and Fig.3. The other insight observed is that increasing the current density does not inhibit the formation of zinc carbonate which can be attributed
Table 1 Experimental results obtained from electrochemical reduction of CO2 at SS 304 L in ACN-0.2 M TBAP in ambient pressure. Exp No.
Current density (mA/cm2)
Cell Potential (V)
Duration of electrolysis (h)
Weight of product (g)
Faradaic efficiency of oxalate (%)
1 2 3 4
5 10 15 30
2.5 3.0 3.5 4.5
1 1 1 1
0.17 0.32 0.46 0.87
42.61 45.96 51.07 81.1 %
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Table 2 Assignment of FTIR bands observed for solid product obtained from CO2 reduction. Wavenumber (cm−1)
Assignment
Compound
1089, 1320 1363, 1472, 1645, 1767 2015 2964
C-O stretch COO− stretch C-C stretch CO32− C=O COO− stretch CN stretch C-H stretch
Zinc carbonate Zinc oxalate Zinc oxalate Zinc carbonate Zinc oxalate Zinc glycolate Zinc cyanide Zinc oxalate dihydrate (adsorbed water molecules) Zinc oxalate dihydrate
1097 1388 1475 1579
3392, 3408
O-H stretch
to the intermediate CO2− anion reacting with CO2 forming CO and CO32− as discussed earlier, See Eqs. (2) and (3) [3]. However, recent studies have found that the formation of carbonate is due to the presence of trace amounts of water in the electrolyte solution. Figueiredo et al. investigated the effect of water content in ACN on CO2 reduction at copper electrodes and reported that the presence of water results in the formation of carbonate and bicarbonate [29]. A more significant finding is that the formation of CO and CO32− happens independently for this system. Hence, the formation of zinc carbonate may be due to trace amounts of water present and the reactions taking place for its formation are shown in Eqs. (4) to (8).
Zn →
Zn2 +
2H2 O +
2e−
+
Based on these results, an optimum current density of 15 mA/cm2 was chosen as there was no formation of byproducts such as cyanides and glycolate for the large scale production of oxalate in a batch reactor operated at 2 bar pressure. 3.2. Electroreduction of CO2 to oxalate in the batch reactor
(4)
2e−
→ H2 +
Fig. 4. Variation of cell voltage during the electrochemical reduction of CO2 to oxalate in the batch reactor at a constant current density of 15 mA/cm2.
2OH−
The schematic of the batch reactor is shown in Fig. 1. It contains provisions for inserting two electrodes separated by a distance of 5 cm and an agitator placed between them. After setting up the electrodes and feeding the electrolyte solution into the reactor, care was taken to ensure the reactor was airtight and CO2 was fed into the reactor for 30 min with continuous stirring for attaining saturation. A constant pressure of 2 bar was maintained inside the reactor. After 30 min, electrolysis was performed by maintaining a constant current density of 15 mA/cm2 and continued for upto 8 h. As the product formation started and continued for longer hours, its accumulation near the electrodes resulted in a reduced area available for CO2 reduction and an increase in the cell voltage. Fig. 4 shows the increase in cell voltage with time when a constant current is maintained. This can be explained due to the fact that the product accumulation near the vicinity of the cathode shields the active sites for the CO2 molecules to be reduced. Further, the accumulated product reduces the migration of electrons coming from anode leading to an increase in resistance of the electrolyte solution. In order to obtain deeper insight into the increase in cell voltage during electrolysis, electrochemical impedance spectroscopy (EIS) was performed in a single compartment cell using SS 304 L as working electrode (area of 2 cm2), Zn as counter electrode and Ag wire as reference electrode. The volume of electrolyte used was 50 mL. Electrolysis was performed at a constant current density of 15 mA/cm2 for 1 h with constant bubbling of CO2. After 1 h, electrolysis was stopped and EIS was performed with respect to open circuit potential (OCP) with frequency ranging from 1 MHz to 100 mHz. This process was repeated until 6 h of electrolysis. An equivalent circuit was fitted using Randles circuit consisting of an electrolyte solution resistance (Rs) in series with a parallel combination of charge transfer resistance (Rct) and a double layer capacitance (Cdl). The obtained values are reported in Table 4. Nyquist plots obtained after every 1 h of electrolysis is shown in Fig.5 (a). The comparison of EIS spectra clearly show an increase in both Rs and Rct with time. The obtained Rs and Rct values were plotted against time of electrolysis and is shown in Fig.5(b). It can be seen from Table 4 that both Rs and Rct values before electrolysis (t = 0) is higher than those obtained after 1 h of electrolysis. This is because once electrolysis starts, CO2 reduction starts facilitating movement of C2O42− and Zn2+
(5)
OH− + CO2 → HCO3−
(6)
HCO3− + OH− → CO32 − + H2 O
(7)
CO32 − + Zn2 + → ZnCO3
(8)
It is therefore crucial to reduce the water content in the electrolyte to prevent the formation of carbonate and obtain oxalate. The removal of trace amounts of water is done using pre electrolysis using two platinum electrodes and applying a constant potential of -1.8 V vs Ag/ AgCl. This removes water present by liberating it as H2 and O2 gas. However, this is a laborious task for the production of oxalate on a large scale as we are dealing with a higher volume of electrolyte solution, typically about 1 L. On the contrary, the use of molecular sieves to remove moisture present in the electrolyte solution can be a better solution and hence, in our studies ACN with 0.2 M TBAP was stored in molecular sieves for 24 h before each experiment. The precise water content in the electrolyte solutions was measured using Karl Fischer titration with an Metrohm 381 K F Coulometer and is reported in Table 3. It can be seen that the amount of water present in 99.9% ACN with 0.2 M TBAP is 623.4 ppm. When this solution is stored in molecular sieves, the water concentration reduces drastically to 37.6 ppm. Although, the formation of small quantities of zinc carbonate could not be avoided, it does not pose a problem during the extraction of oxalic acid from zinc oxalate because zinc carbonate can be removed as salts of concentrated acids which is discussed later in the ester hydrolysis process. Moreover, for longer duration of electrolysis carried in a batch reactor for 8 h, formation of zinc oxalate becomes dominant once all the residual water gets used up in the formation of carbonate. Table 3 Water content of the electrolyte solutions used in the experiments. Solution
Water concentration (ppm)
ACN (99.9 %) + 0.2 M TBAP ACN (99.9 %) + 0.2 M TBAP after removing water using molecular sieves
623.4 37.6
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microscopy and the results obtained are as shown in Fig. 6. It can be seen clearly from these images that there is no significant change in the morphology of the catalyst. The composition characterized by EDX results also reveal no significant change confirming no significant degradation of the elements present in the alloy (Ni and Cr).
Table 4 Solution resistance and charge transfer resistance after various electrolysis duration carried out a constant current density of 15 mA/cm2. Time (hrs)
Solution resistance, Rs (ohm. cm2)
Charge transfer resistance, Rct (ohm. cm2)
Cdl (F/cm2)
0 1 2 3 4 5 6
5.74 4.84 5.20 5.60 6.22 7.13 8.99
4.23 × 104 1.54 × 104 2.28 × 104 2.67 × 104 3.41 × 104 4.00 × 104 5.33 × 104
3.67 × 10−5 3.96 × 10−5 3.29 × 10−5 3.29 × 10−5 2.70 × 10−5 2.31 × 10−5 1.67 × 10−5
3.4. Rate of dissolution of sacrificial anode Another important parameter that has to be considered in this process is the rate of dissolution of the sacrificial zinc anode in the reactor in order to estimate the amount of zinc required for the 8 -h electrolysis. To determine this, the weight of zinc before and after the reaction was measured in each experiment. The initial weight of zinc in the first experiment was 100.13 g and after performing electrolysis for 8 h, the weight of zinc was found to be 82.03 g implying that 18.1 g of zinc was consumed in the reaction. This experimental dissolution rate of zinc was compared with the theoretically predicted dissolution rate of zinc. Pure zinc has a theoretical maximum capacity of 820 mA h/g which is also termed as anode capacity [32]. For 8 h of electrolysis operating at a constant current of 1.8 A (current density of 15 mA/cm2) the dissolution rate was calculated and found to be 17.56 g (See Eq. (11)). This means that theoretically, a minimum of 17.56 g of zinc is required to use it as a sacrificial anode which can continuously supply 1.8 A of current for 8 h. The difference of 0.54 g (18.10 -17.56) obtained in the experiment is due to the zinc operating at around 95% efficiency which means a slightly higher mass of 18.1 g of zinc is required to perform this electrolysis.
ions resulting in a drop in the solution resistance (increase in conductance). However, EIS performed after 1 h of electrolysis show an increase in both Rs and Rct and it increases steadily till 6 h of electrolysis. This is in accordance with the increase in cell voltage obtained in the 8 h of electrolysis performed in the batch reactor at a constant current density of 15 mA/cm2. The reason for this steady rise is essentially due to the accumulation of product (zinc oxalate) formed that blocks the active sites of the catalyst (SS 304 L). Hence, due to an increase in the overall resistance of the system, the cell voltage rises (at constant current density) in accordance with Ohm’s law (V = IR). This phenomenon is similar to potassium carbonate precipitation observed on the cathode side in CO2 electrolyzers using aqueous KOH solution as electrolyte. In such an electrolyzer having an anion exchange membrane, a similar rise in cell voltage is observed due to the precipitated carbonate blocking the active sites of the catalyst [30,31]. It is evident from Fig. 4 that the cell voltage increases steeply after 4 h and reaches a value of 6.7 V at 8 h when product accumulation in the reactor goes very high. This increase in the cell voltage can be overcome by creating a continuous electrochemical process where the electrolyte solution can be allowed to flow through a filtration unit to separate the solid product obtained in the reactor, subsequently recycling the product free filtrate (electrolyte solution) back to the reactor with no/very little solid product in the electrolyte solution. Current efforts are directed towards this task.
Q = I × t = (1.8A × 8 × 3600) = 51, 840 C C = 820mAhg −1 = (820 × 10−3 × 3600) Cg −1 = 2952 Cg −1
w=
51840 C = 17.56g 2952 Cg −1
(9) (10)
(11)
Here, ‘Q’ is the amount of charge applied, ‘C’ is the anode capacity of zinc and ‘w’ is the weight of zinc dissolved. This dissolution rate matched well with the theoretically predicted dissolution rate of zinc. In order to safely operate the electrochemical reduction process, the amount of zinc was kept very high (∼ 100 g) with respect to the minimum 18 g requirement. The total mass of the conversion product obtained from the reactor was found to be 38.1 g. This means that the remaining 20 g (38.1 g18.1 g) comes from oxalate and carbonate ions as the product is a mixture of zinc oxalate and zinc carbonate. This is as expected and importantly no other by-products such as zinc cyanide or glycolate is observed at this current density as discussed earlier. The obtained product was then filtered, washed using ACN, dried and qualitative analysis was done using FTIR spectroscopy which confirmed the presence of zinc oxalate and some amount of zinc carbonate. (See Fig. S2,
3.3. Stability of catalyst The stability of catalyst used is one of the important factors that has to be taken into consideration in electroreduction of CO2. From our studies, the catalyst used (SS 304 L) was relatively stable for 8 h of electrolysis. The increase in cell voltage during electrolysis was due to the accumulation of the solid product in the vicinity of the cathode as explained before and not essentially due to catalyst deactivation. To investigate this, the surface morphology of SS 304 L electrode before and after electrolysis was studied using scanning electron
Fig. 5. (a) EIS spectra of SS 304 L in CO2 saturated ACN-TBAP at open circuit potential after various electrolysis duration. (b) Variation of solution resistance and charge transfer resistance after various electrolysis duration, carried out a constant current density of 15 mA/cm2. 109
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Fig. 6. Scanning electron microscopy images and EDX results of SS 304 L before (a), (b) and after (c), (d) the reaction showing no significant change in the morphology.
Table 5 Experimental results of CO2 electroreduction in the batch reactor with S.S 304 L cathode and Zn anode. Exp No.
Cathode
Current density (mA/cm2)
Cell Voltage (V)
1 2 3 4 5
S.S S.S S.S S.S S.S
15 15 15 15 15
3.5 3.5 3.5 3.5 3.5
304 L 304 L 304 L 304 L 404L
– – – –
6.7 6.6 6.1 7.1
Duration of Electrolysis (h)
Weight of crude product (g)
Faradaic efficiency of oxalate (%)
8 8 8 8 4
38.1 41.2 37.2 35.1 22.4
69.1 66.2 81.1 79.3 77.3
Table 6 Experimental results of the extraction of oxalic acid from zinc oxalate. Exp. No
Weight of zinc oxalate (g)
Volume of ethanol (mL)
Volume of conc. H2SO4 (mL)
Volume of water (mL)
Theoretical weight of oxalic acid (g)
Weight of oxalic acid obtained (g)
Yield (%)
1 2 3 4
2 2 4 4
10 10 20 20
4 4 7 7
15 15 30 30
2.30 2.30 1.15 1.15
0.61 0.65 1.34 1.31
52.9 56.5 58.2 56.9
Supplementary data). The electrolysis experiments for a duration of 8 h were repeated four times and the average weight of solid product obtained was found to be 37.9 ( ± 3.3) g with a Faradaic efficiency of 73.9% ( ± 7.3) for oxalate. The results obtained are shown in Table 5. This increase in the Faradaic efficiency from 51.1% at ambient pressure to 73.9% at 2 bar can be attributed to the increased solubility of carbon dioxide at higher pressure and reduced water content in the electrolyte. The product was then used as the starting material for the synthesis of oxalic acid (See Table 6).
All the marked diffraction peak positions are in good agreement with the standard JCPDS card (Joint committee on powder Diffraction standard) No. 25-1029 confirming that the product is zinc oxalate dihydrate. Both diffractograms show well defined singlet at 2θ = 18.70 for (202) plane of zinc oxalate. X- ray diffraction peak positions of zinc oxalate reported by Guo et.al for the synthesis of rod like zinc oxalate also confirm the same [33].
3.5. X- ray diffraction studies of the obtained product
Fig. 8 shows the XPS of ZnC2O4-C and ZnC2O4-E with high resolution scans for Zn 2p, C 1s, O 1s, peaks. The B.E. value at 1045.1 eV and 1046.1 eV for of Zn 2p1/2 orbital and Zn 2p3/2 orbital is at 1022.5 eV and 1023.5 eV for ZnC2O4-C and ZnC2O4-E respectively [34]. The small shift in B.E values of Zn may be due to difference in the morphology of zinc oxalate as confirmed by FESEM images of both ZnC2O4-C and ZnC2O4-E. Moreover, extra peaks at 1020.8 eV and 1044.1 eV are observed for ZnC2O4-E. This could be attributed to the formation of traces of zinc carbonate in addition to zinc oxalate which is also confirmed by
3.6. XPS analysis of zinc oxalate
In order to confirm the product obtained from electroreduction of CO2 in the batch reactor to be zinc oxalate, X- ray diffraction analysis was performed in addition to FTIR spectroscopy and was compared with ZnC2O4-C. The diffraction pattern obtained is shown in Fig. 7. The XRD patterns acquired clearly show the crystalline structure of zinc oxalate confirming the product obtained in addition to FTIR spectroscopy results. 110
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Fig. 7. X-ray diffraction patterns of (a) ZnC2O4-E obtained at 15 mA/cm2 and (b) ZnC2O4-C.
observed from Fig. 9(a) that ZnC2O4-C composed of micro sheets of different sizes. ZnC2O4-E on the other hand composed of an aggregate of particles, which is evident from Fig. 9(c) and (d). It is also clearly evident from these images that the average particle size of ZnC2O4-C is higher than that of ZnC2O4-E. In order to obtain a clear morphology of ZnC2O4-E and determine its particle size, transmission electron microscopy (TEM) was performed and results obtained are shown in Fig.10. It can be seen that nanosheets of zinc oxalate are formed, having an average length and breadth of 85.2 nm and 32.4 nm respectively. In contrast, the average length and breadth of microsheets of ZnC2O4-C was 461.7 nm and 261.8 nm respectively. This shows that zinc oxalate synthesized from electrochemical reduction of CO2 produces nanoparticles which may be of interest in other applications such as in the synthesis of zinc oxide [37,38]. Further, the thickness of zinc oxalate
FTIR spectroscopy as discussed earlier. Similarly, the high resolution C 1s peaks are qualitatively similar for both ZnC2O4-C and ZnC2O4-C. C 1s peaks for ZnC2O4-C are observed at 288.62 eV and 284.98 eV which can be attributed to the presence of C = O and C-C respectively both of which are present in zinc oxalate and is in close agreement with previously reported results [35,36]. Further, the high resolution O 1s peak is observed at 531.81 eV for ZnC2O4-C and at 531.65 eV for ZnC2O4-E. This small difference in the B.E is possibly due to the traces of zinc carbonate present in addition to zinc oxalate in ZnC2O4-E.
3.7. Morphological studies of zinc oxalate The morphological study of ZnC2O4-C and ZnC2O4-E was carried out with field emission scanning electron microscopy (FESEM). It can be
Fig. 8. High resolution XPS spectra of (i) Zn 2p, (ii) C 1s and (iii) O 1s in (a) ZnC2O4-C and (b) ZnC2O4-E. 111
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Fig. 9. FESEM images of (a) and (b) ZnC2O4-C, (c) and (d) ZnC2O4-E showing a clear difference in their morphology.
Diethyl ether for instance is immiscible in water and forms an organic phase but is not an effective solvent for this extraction process due to the poor solubility of oxalic acid in ethers [39]. In our experiments, calculated amounts of ZnC2O4-E and an excess of 4 N sulphuric acid were taken and mixed vigorously till it completely dissolved. To this solution, calculated amounts of diethyl ether was added and agitated vigorously in a separating funnel which formed two immiscible phases. The organic phase was separated and overnight recrystallization produced crystals of oxalic acid. However, the yield obtained was very low. 2 g of zinc oxalate yielded 35 mg of oxalic acid, which is only about 3.5%. This can be attributed due to the poor solubility of oxalic acid in ethers. In organic solvents such as ethanol and methanol, oxalic acid has a relatively higher solubility [40]. However, these solvents are miscible in water and form azeotropes which makes it difficult to separate it from zinc oxalate.
nanosheets synthesized from CO2 have been probed by atomic force microscopy measurements on mica substrate. The average thickness (Z) was found to be 27.5 nm as depicted in Fig. 11. The average size (X) is in the order of 90 nm (1.1 μm − 0.2 μm) and is in close agreement with the length of nanosheet measured from TEM which is 85.2 nm as mentioned earlier.
3.8. Oxalic acid extraction from zinc oxalate Zinc oxalate obtained from electrolysis experiments was washed, dried and used as the starting material for the extraction of oxalic acid. Zinc oxalate is insoluble in many solvents and dissolves only in strong acids like hydrochloric acid and sulphuric acid. When it is treated with 4 N sulphuric acid, it produces zinc sulphate and oxalic acid (See Eq. (12)). However, both zinc sulphate and oxalic acid are soluble in water. Hence, the use of an organic solvent that is immiscible in water and having a high solubility of oxalic acid in it is essential. It is important that the organic solvent satisfies both these conditions for the effective extraction of oxalic acid.
ZnC2 O4 + H2 SO4 → ZnSO4 + H2 C2 O4
(12)
It is therefore important to separate zinc initially from zinc oxalate since both zinc sulphate and oxalic acid are soluble in water. In the second extraction process that we studied, zinc was first removed from
Fig. 10. TEM images ZnC2O4-E at (a) 200 nm and (b) 100 nm showing the nanosheet morphology of zinc oxalate synthesized from CO2. 112
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Fig. 11. The AFM measured thickness of zinc oxalate nanosheet is 27.5 nm. AFM image is shown with a zinc oxalate nanosheet.
Fig. 12. Flowchart of Ester hydrolysis process used for the extraction of oxalic acid from zinc oxalate.
It is important to note here that some amount of zinc carbonate was also present in addition to zinc oxalate obtained from CO2. However, upon treatment with concentrated sulphuric acid, the zinc carbonate went off as zinc sulphate, releasing carbon dioxide and water which is shown in Eq. (13). The limitation in this process is the release of CO2 and hence its concentration in the product has to be minimized.
zinc oxalate as zinc sulphate using concentrated sulphuric acid (96%) without adding any water. Ethanol with a high solubility of oxalic acid was used as the solvent for oxalic acid extraction. Upon separating zinc sulphate, water was added to the filtrate for performing ester hydrolysis. The schematic of the steps used in this extraction process is shown in Fig. 12. Briefly, ZnC2O4-E was first washed, dried and dispersed in ethanol in calculated proportion and stirred vigorously for 1 h. To this mixture, concentrated sulphuric acid was added and stirred for about 3 h. The white slurry obtained was then centrifuged and the precipitate was separated out which was confirmed to be zinc sulphate. Excess of water was then added to the filtrate solution containing diethyl ester and heated at 700 C for 30 min for the completion of ester hydrolysis. The vapours coming from this reaction were condensed to obtain ethanol which can be reused. Thereafter, the solution was cooled and crystallization was performed multiple times. This yielded white crystals which was confirmed to be pure oxalic acid using 13C NMR spectroscopy.
ZnCO3 + H2 SO4 → ZnSO4 + CO2 + H2 O 13
(13)
C NMR spectroscopy was used for the characterization of oxalic acid crystals extracted from zinc oxalate and was compared with commercially available oxalic acid. For NMR studies, 50 mg of oxalic acid was dissolved in 1 mL of deuterated water (D2O) and the obtained results are as shown in Fig. 13. It can be seen that a single peak at 160.41 ppm is obtained for the extracted oxalic acid and at 161.69 ppm for the commercially obtained oxalic acid showing that the extracted oxalic acid is as pure as commercially obtained oxalic acid. A maximum yield of 58.2% was obtained which is significantly higher than the other process used for separation using diethyl ether. We are currently 113
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Fig. 13. 13C NMR spectroscopy results of oxalic acid, (a) commercially obtained and. (b) extracted from zinc oxalate.
working on methods to improve the yield.
loss encountered in this process.
3.9. Electrolyte and solvent recovery after electrolysis
Declaration of Competing Interest There are no potential conflicts of interest reported by the authors.
One of the main limitations of using non-aqueous solvents for CO2 reduction is its high cost. It is therefore important to recover both the solvent and electrolyte for reuse after the reaction. To study this, the liquid electrolyte used in the reaction was first filtered to separate the solid product. After separation, this liquid was taken in a round bottom flask and loaded into a rotary evaporator. Upon distillation of the solvent ACN, white solid crystals precipitated which was separated and dried. This white crystal was confirmed to be tetrabutyl ammonium perchlorate (TBAP) using FTIR spectroscopy (See Fig.S3, Supplementary data) and a yield of 86.7% was obtained. Both ACN and TBAP can be recovered which becomes an important factor when the economics of this process is considered.
Acknowledgement This work was supported by CSIR Mission mode project on “Catalysis for Sustainable Development (CSD)”, Project No: HCP-0009. The authors are grateful to Central Instrumentation Facility, CSIRCECRI for providing access to FTIR, XRD, SEM and TEM facilities. The authors also thank Mr. Amuthan Dekshinamoorthy for helping us in AFM measurements and Mr. Vitalis Chukwuike for valuble discussions on electrochemical impedance spectroscopy. Appendix A. Supplementary data
4. Conclusions
Supplementary material related to this article can be found, in the online version, at https://doi.org/10.1016/j.jcou.2019.10.011.
In summary, the electrochemical reduction of CO2 to oxalate at SS 304 L cathode in ACN is an efficient method for the production of oxalic acid. Minimizing the water content in the electrolyte solution, operating at an optimum current density of 15 mA/cm2 and maintaining a pressure of 2 bar are the three important factors for obtaining a high yield of zinc oxalate. As the current density is increased to 30 mA/cm2, byproducts like zinc cyanide and glycolate starts to form due to ACN decomposition and partial reduction of oxalate respectively. Although, the formation of small amounts of zinc carbonate cannot be inhibited completely, it does not pose an issue in the oxalic acid extraction process as it can be removed as the salt of concentrated acid. The limitations of this process are the relatively lower yield of oxalic acid (58.1%), release of traces of CO2 during oxalic acid extraction that occurs due to conversion of zinc carbonate to zinc sulphate and the accumulation of product in the reactor that causes an increase in the cell voltage. Average Faradaic efficiency of 73.9% can be obtained for the conversion of CO2 to oxalate and there is a significant increase in the efficiency upon increasing the pressure. ACN and TBAP used in the reaction can be recovered and a recovery rate of 86.1% of TBAP with can be achieved without any byproducts. This becomes a significant factor in process economics due to its high cost. Current efforts are directed towards increasing the yield of oxalic acid extracted from zinc oxalate and designing a continuous process for minimizing the ohmic
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