Article pubs.acs.org/EF
Chlor-syngas: Coupling of Electrochemical Technologies for Production of Commodity Chemicals Tedd E. Lister* and Eric J. Dufek Idaho National Laboratory (INL), Post Office Box 1625, Idaho Falls, Idaho 83404, United States ABSTRACT: This paper describes a novel electrolysis process called chlor-syngas, where synthesis gas is produced at the cathode and chlorine gas is produced at the anode. The work presented is an extension of previous electrolysis system development, where syngas was cathodically generated from water, CO2, and electricity. The process described here uses chloride-based electrolytes. Using HCl as the anolyte provides a low-cost source of Cl−, and leakage of excess protons lowers the catholyte pH, preventing carbonate buildup in the catholyte. Initial electrolysis data are presented here to demonstrate the feasibility of the process in KCl/KCl and KCl/HCl electrolytes. Using the electrolysis data, an estimation of the energetic and environmental benefits is presented. The process could be a path to a more sustainable chemical industry, where the starting materials are low-value or wastes from other related processes.
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INTRODUCTION Efficient chemical processing schemes use readily available materials and produce useful products with minimal waste. The process described here, chlor-syngas, produces two commodity gas streams, Cl2 and synthesis gas (syngas), using low-value chemicals, CO2 and HCl. The two product streams account for a vast array of chemical products, which could lead to a sustainable plastics industry and potentially lower atmospheric CO2 (release or capture). The chlor-syngas process could replace two existing processes in current use: (1) chlor-alkali for the production of Cl2 and (2) gasification of fossil sources, such as natural gas or coal, to produce syngas. The proposed scheme is analogous to chlor-alkali with two major differences: (1) the anode uses recycled hydrochloric acid (HCl) recovered from organic chlorination processes in place of alkali metal chlorides, and (2) the cathode produces syngas instead of H2 and caustic. This configuration dovetails well into significant chemical processes and is designed to be efficient, clean, and sustainable. A significant amount of work has been performed on the CO2 reduction half reaction, where products, such as CO,1,2 methane and ethylene,3,4 formic acid,5,6 and alcohols,7 are produced. A key factor in determining the product distribution has been the cathode material.8−10 Pt group metals tend to favor hydrogen evolution because of their strong adsorption of CO and hydrogen. Metals, such as Ag, Au, Zn, and Cu, have weak CO adsorption and moderate hydrogen overvoltage, leading primarily to the formation of CO, although Cu electrodes form a host of hydrocarbons. Formic acid is the primary product for poorly catalytic metals (Hg, Cd, Pb, In, and Sn), because C−O bonds are not broken without surface adsorption. Because of limited CO2 solubility in aqueous solutions, the turnover rate for CO2 reduction is masstransport-limited.11 Gas diffusion electrodes (GDEs), which are porous electrodes with the reaction gas fed from the back side of the electrode, have been used to increase the effective current densities.4 A limited number of studies have reported on scale-up ready electrolysis systems for CO2 reduction.2,12,13 Previous work in © 2013 American Chemical Society
this lab has demonstrated commercial current densities for syngas production by employing a GDE.14 A commercial GDE material called Silflon was used because of the demonstrated stability as an oxygen depolarizing electrode (ODE).15 GDEs possess significant porosity, with the catalytic material typically held in a fluoropolymer matrix. Electron transport occurs at electrode−gas−electrolyte boundaries (three-phase boundaries) within the GDE.16 Higher current densities can be realized by rapid gas replenishment at the three-phase boundaries within the electrode. Industrial electrochemical production of syngas may require use of economical sources of CO2. This increases the likelihood of exposing electrode materials to contaminants, including sulfur compounds, which exist in some sources at levels up to 0.1 ppm.17 Sulfur and S-containing organic species are wellknown poisons for many systems because of the strong adsorption to a wide variety of metal surfaces, including Ag. Deliberate poisoning of S on Ag GDEs showed some decrease in reactivity to CO2; however, the effect is mostly reversible.18 Poisoning with S resulted in a depolarization of the H2 evolution reaction, as evidenced by lowered cathode potentials and high Faradaic efficiency for H2 over CO. It was found that welding-grade CO2 could be used with minimal loss in CO2 conversion. It is thought that, because of the negative potentials at the cathode during electrolysis, S species are reduced to H2S and liberated. Chlor-alkali is the established electrochemical process used to produce Cl2 and alkali (metal hydroxide).19 Specifically, the reactions are (1) oxidation of chloride to Cl2 at the anode and (2) reduction of water to H2 gas. Alkali is produced by consumption of protons during water splitting. This industry has seen use of several different electrolyzer types over the Special Issue: Accelerating Fossil Energy Technology Development through Integrated Computation and Experiment Received: December 10, 2012 Revised: January 16, 2013 Published: January 18, 2013 4244
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exchange membrane to prevent product gases from the anode and cathode from mixing. The process should be a net generator of water, which provides additional benefits. Chlor-syngas dovetails into significant chemical processes because of the synergy of the reactant and product streams. Figure 2 describes the relationship of chlor-syngas to other related processes. The process uses two low-value feedstocks: CO2 and HCl. The product streams account for a vast array of chemical products. The process is driven by electrical energy, which can be obtained from clean sources, such as wind, solar, hydro, or nuclear. In addition to reusing HCl, the process also serves to either sequester CO2 into organic precursors, which can produce commercial products, such as polymers, or generate carbon-neutral liquid fuels, which can be directly used within the existing U.S. transportation infrastructure. CO2 is a highly dilute gas in the earth’s atmosphere.20 Processes are being developed to efficiently concentrate CO2 from the atmosphere by direct air capture (DAC).21 More concentrated sources are also available from industrial processes. Industrial practice employs cold methanol or methyl diethanol amine (MDEA) to physically adsorb CO2, as used to condition syngas derived from gasification.22,23 Concentrated sources of CO2 are available from coal-fired power plants.24 Conventional syngas generation liberates CO2, which could be converted to reduce emissions. Carbon capture and storage (CCS) proposes to permanently store CO2. Alternatively, beneficial use considers CO2 as a feedstock to other industrial processes, such as what is proposed here. Hydrochloric acid (HCl) is a byproduct of many organic reactions involving Cl2 and phosgene.25 These reactions are an important part of a number of chemical products, such as plastics, rubber, and others. These reactions produce HCl as a byproduct that can be sold as a commodity chemical but is produced beyond the consumption rate and, thus, has a low value.25 Methods to recycle HCl back into the process stream are desirable.
years. Mercury cells (Castner−Kellner cell 1892) employ a liquid mercury pool as the cathode, where a Na−Hg amalgam is formed. The pool travels by gravity in a trough over the length of the cell and then is stripped of sodium in a denuder, which creates a concentrated sodium hydroxide solution. The anode reaction is performed on dimensionally stable anodes (DSAs), which are titanium with catalytic coatings containing Ru or Ir oxides. Diaphragm cells (Griesheim cell 1885) are named for the physical barrier used to separate the solutions. These cells use a more conventional electrolyzer design, with final products formed directly at the electrodes. Brine (NaCl or KCl) is fed into the anode compartment, where Cl2 is generated at the DSA. Hydrogen and caustic are generated at the cathode. The diaphragm, often made from asbestos in older systems, allows brine to diffuse through the diaphragm to balance charge while preventing free mixing of solutions. Specifically, the separation of Cl2 and caustic that form hypochlorite is achieved. Finally, membrane cells have been employed more recently (since the early 1970s) to more efficiently separate the electrode compartments. Developments include reducing O2 in a GDE in place of water reduction to increase efficiency through a drop in the cell potential.16 Electrolysis processes perform separate reactions at the anode and cathode. To design a commercially viable process, these reactions should both produce an economically useful product. Figure 1 describes the chlor-syngas process. The
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EXPERIMENTAL SECTION
Experiments were performed using a commercially available filter press cell, which was modified in-house to accommodate a GDE (model AB, ElectroCell Systems). The anode (Ru-based dimensionally stable anode, ElectroCell Systems) and cathode (Ag GDE, Silflon, Gaskatel) had active geometric areas of 10 cm2. All data presented in this paper were performed using the same electrode to demonstrate durability and reactivity. Experiments were performed using a Nafion 115 cationexchange membrane with a thickness of 0.127 mm. High-purity CO2 (Norco) was passed through a hydrocarbon filter (Restek) and delivered to the back of the GDE at controlled flow rates using a massflow controller (1479A, MKS, Inc.). All experiments described here used a 30 mL/min flow rate. Solutions were prepared using 18.2 MΩ cm water and ACS-grade or better chemicals. Cathode potentials (Ec) are reported versus a Ag/AgCl reference (3 M NaCl, Bioanalytical Systems), which was maintained at room temperature (22 ± 2 °C) in an external compartment connected through a salt bridge. Experiments were performed using a constant current density of 75 mA/cm2 provided by a 3 A and 30 V direct current (DC) power supply (BK Precision). The cell potential (Ecell), Ec, and CO2 flow were captured using an in-house LabVIEW program. Unless specified, electrolysis experiments were performed at room temperature (22 ± 2 °C). The anolyte and catholyte volumes were 100 mL. A multi-channel peristaltic pump cycled the electrolyte between the cell and electrolyte reservoirs at 15 mL/min. C-flex tubing (Cole Parmer) was used to connect the electrolyzer to the reservoir and pumps. This tubing was chosen for chemical compatibility and low gas permeability. Some swelling was observed in the anolyte tubing
Figure 1. Diagram of the chlor-syngas electrochemical process.
electrochemical reactions involved in the chlor-syngas process are included in Figure 2. The total cell reaction assumes a 1:1 ratio of H2 and CO products. The electrolyzer uses a proton-
Figure 2. Diagram showing how chlor-syngas would fit into commodity chemical production. 4245
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exposed to Cl2, however. The catholyte pH was monitored using a flow-through pH probe (Microelectrodes, Inc.) located between the exit of the cell and the catholyte reservoir. Product analysis was conducted using a previously described protocol on either a Hewlett-Packard 5890 gas chromatograph (GC) or a model 6890 GC (Hewlett-Packard).14,18 The cathode gas was flowed directly to the GC gas sampling loops. The syngas ratio was reported by dividing the compositions of H2 to CO as determined by GC. Product gas generation was monitored with a digital flow meter (Agilent ADM-1000). From GC analysis, the syngas ratio was determined. CO and H2 were the only reduction products detected. Chlorine gas was captured in a two-trap system, where the headspace of the anolyte exited by tubing into the aqueous phase of a secondary vessel containing 10 g of KI in 200 mL of deionized water. The reaction produces iodine, which is reddish brown in color and alleviates chlorine gas venting. The secondary vessel headspace exited into the solution of a second smaller vessel with 1 g of KI in 20 mL. This setup was effective, because the secondary container did not turn color in any experiment. At the end of each experiment, the anolyte solution itself was also reacted with KI to capture dissolved Cl2 or ClO−. The iodine-containing solutions were combined and titrated using a 1 M solution of K2S2O3 until iodine color vanished from solution. This method likely underestimates Cl2 gas generated but provides a useful estimate.
Figure 3. Electrolysis polarization curves and syngas ratio in 2.5 M KCl. The syngas ratio was determined after 20 min at each applied current.
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RESULTS AND DISCUSSION While most reports have used carbonate-based electrolytes to perform CO2 reduction, previous work has demonstrated CO2 reduction to CO in KCl.26 This work, which focused on the CO2 reduction half reaction, demonstrated that CO could be produced at high Faradaic efficiency above pH 2.5. Below a pH of 2, the efficiency was found to drop to near zero. Previous work in this laboratory, where total cell optimization has been the focus, used sulfate-based catholytes and KOH-based anolytes. During operation, carbonate salts form in the electrolyte because of CO2 reaction with hydroxide, generated by the basic conditions caused by proton consumption described in reactions in Figure 2. The anode evolves O2, a reaction that consumes hydroxide and, thus, must be constantly replenished using that electrolyte system by either the addition of hydroxide or altering the cell configuration to allow for transport of hydroxide from the cathode.27 This also results in the production of a low-value gas that would likely be vented to the air. Naturally, a more valuable anode product would improve the economic viability of the process. Thus, initial investigations focused on the use of KCl electrolytes, where Cl2 was formed at the anode and syngas was formed at the cathode. The polarization and cathode product behavior is shown in Figure 3. These data were collected by stepping from low to high current values, where a sample of cathode gas was sampled at each current, after 20 min, to determine the syngas ratio (H2/CO). In a fashion similar to what has been observed in sulfate electrolytes, at low current densities, the syngas ratio is low, with CO being the major product. As the current is increased, the production of H2 increases, as observed previously. The syngas ratio measured here is significantly lower than that measured in K2SO4 electrolytes under comparable conditions.14 Syngas ratios for producing chemical or fuel products range from 1 to 3 depending upon the specific chemical and process. While not demonstrated here, the CO2 flow rate applied to the back side of the GDE can also be used to tailor the syngas ratio.14 The electrolysis data in Figure 4 were collected at a constant current of 75 mA/cm2 over 3 h of operation. Note that the potential values are quite stable and the syngas ratio is similar to
Figure 4. Electrolysis data (top) and syngas ratio (bottom) using 2.5 M KCl electrolytes at 75 mA/cm2.
that produced in sulfate electrolyte in the same cell and electrode types. The cell potential (Ecell) value is slightly higher, while the cathode potential (Ec) is similar to previous work.14 Thus, the change in the anode reaction is the likely source of the potential increase. The increase in the syngas ratio appears to be reversible, because the same cathode was used to generate all data shown in this paper (and other data not shown). During the electrolysis, the catholyte pH increases slightly. This is due to the formation of carbonate as protons are consumed in the cathodic reactions. This would present an added complexity and cost to long-term operation of a commercial electrolyzer. However, the generated carbonates could be a useful third product. Experiments were performed at elevated temperature for comparison. Industrial-scale electrolyzers operate at elevated temperature because of resistive and kinetic losses within the cell. As observed previously, the temperature has a beneficial effect on Ecell and Ec.14 As a comparison, at 70 °C, Ecell was 3.53 V and Ec was −1.56 V at 100 mA/cm2. The corresponding room temperature values (from Figure 3) were 4.49 and −1.99 V. The increased temperature appears to have the largest effect on activating H2 evolution as the syngas ratio for this 4246
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comparison increased from 0.32 to 1.32 at 70 °C. The final value is still at the lower range for syngas compositions. Electrolysis Using HCl Anolyte. Experiments were performed using HCl, a low cost and available source of chloride. As described earlier, HCl is produced in organic chlorination reactions and has a low commodity value compared to the sodium or potassium salts, which must be mined. A key concern in using HCl anolyte is excessive transport of protons across the membrane. As shown in previous work, lower pH results in a sharp decline in CO production at planar electrodes.26 However, a balanced excess of protons, which leads to slightly acidic pH values, could eliminate issues with carbonate formation. Using the same cell configuration, the anolyte was replaced with 1 M HCl while retaining 2.5 M KCl in the catholyte. The electrolysis data using this configuration are shown in Figure 5. Over the course of the
expected that the efficiency increases during the experiment. The tubing used was rated “A” in chemical compatibility with wet chlorine but only “B” for bleach liquors. The tubing was chosen for low transport of H2, CO, and CO2, although no value was provided for the tubing. There were signs of tubing swelling on the anode side of the cell, particularly leading to the trap system. This could have led to losses in recovery of Cl2 as well. This issue was not thoroughly investigated because industrial generation of Cl2 approaches 100% current efficiency and competition with O2 evolution is typically not an issue.19 These experiments were tailored around determining the performance of the syngas-producing cathode. The final data point in Figure 5 suggested an advantage of using higher HCl concentrations to perform the chlor-syngas electrolysis. Thus, an experiment was performed using 2 M HCl (Figure 6) to assess operation and if the higher values result in
Figure 6. Electrolysis data (top) and syngas ratio (bottom) at 75 mA/cm2 using 2 M HCl as the anolyte and 2.5 M KCl as the catholyte. Figure 5. Electrolysis data (top) and syngas ratio (bottom) at 75 mA/cm2 using 1 M HCl as the anolyte and 2.5 M KCl as the catholyte. Note that last data point was taken after the addition of 10 mL of concentrated HCl to the anolyte.
decreased CO generation because of low catholyte pH. The higher HCl concentrations resulted in additional acid transport to the cathode, as evidenced by the lower cathode pH during operation compared to Figure 5. However, pH does not appear to affect the production of CO, as described in previous work.26 It is thought that pH is less of an issue for the GDE used in the present study because of the enhanced hydrophobicity provided by the fluoropolymer matrix. Inside the GDE, protons would be consumed because of either cathode reaction. Thus, the localized pH within the electrode is likely higher than in the bulk of the catholyte because of the confined pore space. This presents the prospect of using a high HCl concentration in the anolyte while maintaining activity for CO production. Note that, while the same electrode was used for Figures 4−6, the syngas ratios are observed to increase from less than 0.1 at the start of each experiment. The repeatability of the three independent runs suggests that the increase in the syngas ratio during each run is not the result of a permanent electrode degradation process. One possible explanation is the gradual change in wetting within the GDE as an experiment proceeds. At this time, this appears to be the best explanation for this effect. The Faradaic yield for Cl2 was 76%, slightly higher than that observed in 1 M HCl.
experiment, the pH drops sharply to below a value of 2 and stabilizes at ∼1.3 after approximately 30 min. As described in the reactions in Figure 2, protons are consumed in the catholyte for both H2 and CO generation. Thus, the decrease in pH suggests additional proton transport. The syngas ratio increases slightly over the course of the electrolysis but remains a CO-rich stream throughout. The syngas ratio remains below that observed with KCl as the anolyte (Figure 4). Ec is stable, while Ecell increases in a steady fashion over the electrolysis run. This increase is due to the consumption of Cl−. This was confirmed by the final data point (at 132 min), where a 10 mL aliquot of concentrated HCl was spiked to the anolyte 5 min beforehand. This decreased Ecell to below the value at the start of the experiment. Note that final Ecell (3.54 V) is 0.5 V below that for the initial value in 2.5 M KCl (4.03 V). The chlorine gas captured in the trap was analyzed by titration, giving a Faradaic efficiency of 70% over the entire experiment. It was observed that bubbling of gas to the second trap was greatest at the initiation of the experiment and decreased later. Thus, it is 4247
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A final experiment was performed to determine the potentials that would be achieved by increasing the cell temperature to 70 °C. At 75 mA/cm2, Ecell was 2.91 V, Ec was −1.47 V, and the syngas ratio was 0.76. Increasing the current to 150 mA/cm2 resulted in Ecell of 3.52 V, Ec of −1.76 V, and a syngas ratio of 1.53. Projected Energetic and Environmental Benefits. The data above provide the basis for a preliminary examination of the energetic and environmental benefits of generating Cl2 and syngas using a single electrochemical cell. The energy efficiency, ε, of a process can be defined by the enthalpies of the formed products (at their respective Faradaic efficiencies, η) and Ecell as shown in eq 1 ε=
efficiency as producing Cl2 and syngas independently using current high-efficiency processes. While comparable energy requirements are important, the prime advantage of the chlor-syngas system is the ability to use waste products to generate commercially relevant chemicals. In particular, the generation of syngas using coal or natural gas can be viewed as a net CO2-neutral technology at best if it is assumed that complete capture occurs while the chlor-syngas process uses CO2 to produce syngas. At a syngas ratio of 1.5, 245 kg of CO2 is converted for each ton of Cl2 produced. For reference in 2010, roughly 10 million tons of Cl2 was produced in Europe, with the North American total being closer to 13 million tons.30,31 While the quantities of Cl2 produced on the two continents are similar, the distribution of electrolysis cells varies dramatically, with mercury cells accounting for 30.6% of the European production capacity.30 These cells, which are already scheduled for replacement, are less efficient than membrane cells, and thus, the chlor-syngas cell described here would provide an improvement in energy efficiency. At a 10% entry into the European production capacity for Cl2, 245 000 tons of CO2 could be converted from a present waste to a useful chemical precursor. In North America, only 4% of the production capacity is from mercury cells.31 Replacement of this Cl2-generating capacity would convert 127 000 tons of CO2 annually.
ηΔHCO + ηΔHH2 2FEcell
(1)
where F is Faraday’s constant. This equation was used in previous reports, where the anode produced O2.2 With the anodic product being Cl2, two competing chemical reactions can be written. CO2 + 2H+ + 2Cl− → CO + H 2O + Cl 2 ΔH = 331 kJ/mol 2H+ + 2Cl− → H 2 + Cl 2
(2)
ΔH = 334 kJ/mol
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(3)
CONCLUSION This paper has presented a novel process where syngas and chlorine productions have been integrated into a single process. Because the syngas generation process has low energy efficiency, coupling with a useful anodic process provides additional values and offsets those losses. Initial work showed that syngas production was not adversely affected by the chloride electrolyte. The process could be operated much like previous experiments where sulfate electrolytes were employed. To make use of a lower cost source of Cl−, experiments using HCl as the anolyte were performed. While excess protons leaked across the membrane to the catholyte lowering pH, the syngas ratio was not adversely affected and the lack of carbonate buildup was an advantage. The Faradaic efficiency for Cl2 was lower than expected, which could either be sampling error or the production of O2 in favor of Cl2 at the start of the electrolysis runs. Industrial production of Cl2 does not suffer from significant losses because of O2 production and should not be a future problem. Using the ideal electrolysis data, an estimation of the energetic and environmental benefits is presented. Given adequate development, chlor-syngas cells could operate at cell voltages and production rates similar to membrane cells. The prime environmental benefits are the ability to replace mercury used in aging mercury cells and the consumption of CO2 to generate an important chemical precursor. The process could be a path to a more sustainable chemical industry, where the starting materials are low-value or wastes from other related processes.
By including the anode contribution, at 70 °C, ε is 0.59 at 75 mA/cm2 and 0.49 at 150 mA/cm2. For comparison operation under similar conditions, using O2 generation at the anode results in slightly lower values for ε.14 Currently, Cl2 and syngas are manufactured in separate processes. When the two processes are combined, environmental and energetic benefits could be achieved. Preliminary energy calculations were performed using the data collected above and data currently available for chlor-alkali generation and the production of syngas from coal. The calculations are based on the production of 1 ton of Cl2. At a 1.5 syngas ratio for every ton of Cl2 generated, 0.17 tons of syngas are produced. Commercial membrane cells are generally accepted as being the most energy efficient electrolyzers widely used for chlor-alkali production. Typical operating conditions for these cells range from ∼500 to 1000 mA/cm2 with cell voltages of slightly over 3 V. At 500 mA/cm2 and 3.07 V (at 90 °C), roughly 2300 kWh is required to generate 1 ton of Cl2 gas.28 For comparison, the present data if scaled to the production of 1 ton of Cl2 would require between 2200 kWh (75 mA/cm2) and 2660 kWh (150 mA/cm2). Thus, the chlor-syngas system is capable of operating near the range of membrane electrolyzers that are already in use (albeit at a lower throughput). The chlor-syngas process is at the initial stage of development and, thus, has also not been optimized to the extent of a commercial process. However, unlike the chlor-alkali cells, the second product from a chlor-syngas cell is syngas, which is produced from both coal and natural gas sources. The generation of syngas from coal in the U.S. for the production of chemicals accounted for 3 million tons of the product in 2010, with the average ton of syngas requiring 1.61 MWh to produce.29 When the energy requirements of producing the two products are combined, roughly 2600 kWh is required to generate the same quantities generated using a chlor-syngas cell operating at 150 mA/cm2. Thus, chlor-syngas technology is of comparable energy
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest. 4248
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(28) Bommaraju, T. V. Salt chlor-alkali and related heavy chemicals. Kent and Riegel’s Handbook of Industrial Chemistry and Biotechnology; Springer Verlag: New York, 2007; Vol. 2, Chapter 26. (29) National Energy Technology Laboratory (NETL). 2010 World Gasification Database; http://www.netl.doe.gov/technologies/ coalpower/gasification/worlddatabase/index.html (accessed Nov 27, 2011). (30) Euro Chlor. Chlorine Industry Review 2010−2011; Euro Chlor: Brussels, Belgium, 2011. (31) Chemical Market Associates, Inc. Chlor-alkali Market Report; Chemical Market Associates, Inc.: Houston, TX, Jan 17, 2011; Issue 200.
ACKNOWLEDGMENTS We acknowledge Dr. Lenny Scott of Olin Chemical (Augusta, GA) for providing helpful discussion in the development of this concept. We also acknowledge Simon Stone at Giner, Inc. for helpful discussions about this work. Work was supported through the INL Laboratory Directed Research and Development (LDRD) Program under DOE Idaho Operations Office. This manuscript has been authored by Battelle Energy Alliance, LLC under Contract DE-AC07-05ID14517 with the U.S. Department of Energy.
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