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Energy storage by the electrochemical reduction of CO2 to CO at a porous Au film
Journal of Electroanalytical Chemistry 526 (2002) 125– 133 www.elsevier.com/locate/jelechem
Energy storage by the electrochemical reduction of CO2 to CO at a porous Au film Gregory B. Stevens *, Torsten Reda, Burkhard Raguse CSIRO Di6ision of Telecommunications and Industrial Physics, PO Box 218, Lindfield, NSW 2070, Australia Received 16 August 2001; received in revised form 4 January 2002; accepted 6 January 2002
Abstract We have investigated the performance of a porous Au film electrode for future applications in energy storage by the reduction of CO2 to CO. The electrode consists of a porous hydrophilic polymer membrane with a 260 nm Au film, applied by vapour deposition. The hydrophilic property of the polymer substrate was necessary for obtaining current densities comparable to those at an ordinary Au electrode under similar conditions: aqueous KHCO3, ambient temperature and approximately one atmosphere pressure of CO2. At potentials more negative than the activation potential for CO2 reduction, − 1.2 V versus SCE, the current density was of fractional order with respect to the CO2 pressure, indicating the involvement of an adsorbed reaction intermediate, possibly a radical anion complex (CO2)2−. Analysis by GC and FTIR of the gas reaction product showed that the energy storage efficiency was maximal near the activation potential. Close to this potential, CO was formed with a Faradaic efficiency of 75% and an enthalpic efficiency of 35%. Under these conditions, the power storage capacity of the electrode was 50 W m − 2. © 2002 Elsevier Science B.V. All rights reserved. Keywords: Carbon dioxide; Au electrode; Carbon monoxide; Energy; Electrocatalytic reduction; Current efficiency
1. Introduction As natural oil reserves become depleted, attention will increasingly focus on alternative energy sources. These may include hydrocarbons and oxygenated aliphatics, synthesized from CO and H2 via the Fischer –Tropsch process. The energy and raw materials for this are currently derived from the burning of coal, with the accompanying release of CO2 as a by-product. However, concerns that increases in atmospheric CO2 may lead to serious global climate changes have led a number of authors to suggest that CO2 itself may be used as a source of carbon for the production of petroleum-like materials [1 – 3]. This may then lead to the possibility of regulating the concentration of atmospheric CO2 [4,5]. As CO2 is one of the most thermodynamically stable carbon compounds, a highly energetic reductant or an external source of energy is required to convert it into * Corresponding author. Tel.: +61-29413-7490; fax: + 61-294137202. E-mail address: [email protected] (G.B. Stevens).
other carbon compounds. It is well known that CO2 can be reduced electrochemically [6,7]. The mechanism of this begins with the formation of an adsorbed CO2 radical anion intermediate [8]: CO2 (ads)+e − CO2 − (ads)
(1)
The dominant reaction, subsequent to the initial formation of CO2− (ads), depends on factors such as the type and pH of the electrolyte, the electrode potential, stirring, the CO2 partial pressure and the electrode material [9]. According to Hori et al. [10], in aqueous media transfer of a proton from water to an adsorbed radical anion results in the formation of CO: CO2 − (ads)+H+ + e − CO + OH −
(2)
This is catalysed by metals in the order Au\ Ag \Zn. In aprotic electrolytes, adsorbed radical anions combine with CO2, resulting in the formation of CO and CO23 − by disproportionation [11]: 2CO2 − + e − CO + CO23 −
0022-0728/02/$ - see front matter © 2002 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 2 - 0 7 2 8 ( 0 2 ) 0 0 6 8 8 - 5
(3)
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At other metals from Groups IIB, IIIB and IVB in aqueous electrolytes, formate is the dominant reaction product in the order of activity Cd\ Sn\ In \ Pb \ Tl \Hg [12,13]. This is due to weak binding of CO2 radical anions at these metals. After forming at the electrode, free radical anions are protonated to produce formate ions, whereas in aprotic solvents, free radical anions dimerize to form oxalate [14,10]. In acids, CO2 is reduced to methanol on Mo [15] and to hydrocarbons at Cu electrodes [16]. Other metals show little activity in reducing CO2, either because the hydrogen overpotential is too low or because CO binds tightly to the metal, preventing further reactions with CO2 from taking place. The low solubility of CO2 in aqueous electrolytes is a limiting factor in the reaction rate and, as a result, the reaction will be diffusion limited. The mass transport limitation can be overcome to some extent by the use of gas diffusion electrodes (GDEs), which interface the gas, electrocatalyst and electrolyte phases [1]. Typically, GDEs are constructed from metal-impregnated polytetrafluoroethylene (PTFE)-bonded carbon powders [17], in which the CO2 diffuses through a porous hydrophobic region in the GDE until it reaches the electrolyte in a hydrophilic region, where it is reduced [18]. Although high current densities can be obtained with this type of electrode, the hydrophobic region is prone to flooding. Cook et al. report that this results in a gradual decline of the current density [16]. This problem may be overcome by the use of a solid polymer electrolyte (SPE) such as a perfluorosulfonate ion-exchange resin (Nafion®). Such electrodes have been used in the reduction of CO2 to CH4 and C2H4 [19,20]. However, Nafion®-based electrodes work best in acid media and only relatively small Faradaic efficiencies and current densities have been achieved in the reduction of CO2 with these GDEs. To minimize the adverse effects of mass transport in GDEs, the thickness of the active layer should be kept to a minimum [21] which also leads to economic utiliza-
tion of expensive electrocatalysts. Recently, Chatterjee and Basumalick reported the use of a lightweight GDE made from a thin Ag film supported on a porous PVC membrane, for use in fuel cells [22]. Because this type of GDE can be constructed with an inert substrate, it can be used with electrolytes in which CO2 is more soluble. Although Au is highly selective for CO formation, it has not been studied in great depth for this purpose. Shibata et al. investigated the use of an Au impregnated PTFE-bonded carbon GDE for the electrochemical reduction of CO2 and nitrate to urea [23]. In this work we investigate the electrochemical reduction of CO2 to CO in aqueous media at a thin porous Au GDE and consider the potential application of this system for energy storage.
2. Experimental Throughout this paper, the potential of the working electrode (EW) is measured with respect to a standard calomel electrode (SCE) and current densities are expressed in terms of the geometrical area of the electrode.
2.1. Electrode preparation The film electrodes were prepared by vapour deposition of Au onto one side of porous membranes made of polyvinylidene fluoride polymer (Millipore GVWP), polyethylene (DSM 7P20) and PTFE (Millipore FGLP). Gold (99.99% Au) was evaporated under vacuum (Edwards model 12EA/461), directly onto membranes located 30 cm above an electrically heated tungsten boat. The evaporation rate was 0.2–1.0 nm s − 1 and was continued until the final thickness was 200– 260 nm, as measured with a thickness monitor (Maxtek TM-100). The crystal structure of the metal film was investigated with an X-ray diffraction spectrometer (Philips X’pert). Microscopy and microanalysis was done with a scanning electron microscope (Jeol JSM-6300F) equipped with an energy-dispersive X-ray spectrometer.
2.2. Cell and electrolyte
Fig. 1. Schematic diagram of the electrochemical cell. R, saturated calomel reference electrode; C, Pt wire coil counter electrode; E, 500 mM KHCO3 electrolyte; W, working electrode; G, gold leaf contact to the working electrode; O, ring seal; P, inlet pressure gauge; F, outlet pressure gauge and flow meter.
The film electrodes were supported in an electrolysis cell illustrated in Fig. 1. The cell body was made of borosilicate glass and the gas cavity of methyl methacrylate polymer. Electrical connection to the Au film was made using Au foil. A nitrile ‘O’ ring between the membrane and the cell housing formed a fluid seal, leaving an electrode area of 3.17 cm − 2. Cyclic voltammograms were measured with an electrochemistry recording unit (AD Instruments).
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Gas samples were injected into a Porapak Q column in a gas chromatograph equipped with a thermal conductivity detector (3400 GC Varian Inc.). Gas products were identified by comparison of the retention times with those of a reference gas. Quantitative measurements of CO and CO2 in the gas mixtures were obtained as the ratio of the peak areas of gas reaction products and the same volume of pure gas samples. The effluent gas was also analysed by measuring the absorption in the infra-red region using a Fourier transform infra-red spectrometer (BioRad FTS-65). The absorption integrated from 2200 to 2100 cm − 1 and from 2277 to 2255 cm − 1 was used to measure CO and CO2 concentrations in the gas outflow.
Fig. 2. Scanning tunnelling microscope image of the surface of the porous Au layer. The lighter regions indicate Au. EDX microanalysis did not show the presence of O or C, indicating that Au completely covered the substrate layer.
Measurements of the rate of CO formation were made after passing CO2 over a series of eight Au/ GVWP cathodes, so as to increase the volume of reaction products for analysis by gas chromatography and infra-red absorption. In this case, the membranes were supported in a methyl methacrylate polymer housing. A common counter electrode made of a Pt wire (0.5 mm dia× 540 mm) was located 5 mm from the membranes. Impedance spectra of the eight cell unit were measured using an electrochemical impedance system (EG&G Instruments model 263). The electrolyte was made from 99.99% KHCO3 and MilliQ water. The electrolyte was prevented from flowing into the gas phase by maintaining the gas pressure at 10 kPa above atmospheric pressure. Gasses used were Ar (99.99%) and CO2 (99.95%), which were supplied at pressures of between 10 and 50 kPa and at flow rates of up to 20 ml min − 1 via copper and FEP plastic tubing.
2.3. Electrolysis and analysis of products In the cyclic voltammetry work, traces of CO in the effluent gas were detected with a colorimetric CO detector (Onix Process Analysis). For the quantitative analysis of gas products, a current was passed for periods of up to 3 min at potentials of between − 1.0 and − 3.6 V. The electrolyte temperature was maintained below 26 °C by allowing for cooling between measurements. Gas products from these cells were passed through a 9 cm column of powdered silica gel to exclude water vapour before analysis. The gas flow rates were calculated from the time taken to collect a volume of gas at atmospheric pressure.
3. Results and discussion
3.1. Electrode structure The electrical resistance of the films was measured by positioning two Au bars, each 1 cm long and 2 mm wide, on opposite sides of an imaginary 1 cm square on the film surface. The resistance between the bars on these surfaces was between 1.6 and 2.0 V. However, when used with the KHCO3 electrolyte, Au electrodes made on hydrophobic membranes (DSM 7P20 and FGLP) had much lower current densities than those made on hydrophilic membranes (GVWP). The relatively high current density obtained with the GVWP membranes was attributed to its hydrophilicity. Fig. 2 shows an SEM image of the surface of the porous Au film evaporated onto a GVWP membrane. The surface appears to be sponge-like with pores of varying sizes of up to 1 mm across. The larger pores are much larger than the nominal pore diameter of the substrate layer, 220 nm. This suggests that structural changes to the GVWP membrane occurred during the evaporation, resulting in the Au film bonding well to the GVWP membrane. This was evident when sticky tape was applied to the Au layer and did not remove it when peeled off. The X-ray diffraction spectra of the film indicated the presence of polycrystalline Au, with a preferred (111) orientation and an average crystallite size of about 60 nm. Fig. 3 shows a schematic of the cross section of a membrane pore to indicate the assumed geometry of the gas electrode electrolyte interface. According to Bockris and Cahan’s Finite Contact Angle Meniscus model of porous electrodes [24], most of the current associated with the reduction of the gas phase will be produced very close to the meniscus tip. The electrolyte contacted the gas phase after passing through the membrane by capillary action. As the membrane became wet, it was assumed that the total surface energy of the water and membrane decreased
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and would need to be raised before gas will pass through to the electrolyte. This will occur at pressures above 340 kPa for the GVWP membrane [25].
3.2. Transient current-6oltage characteristics Cyclic voltammograms of the Au electrode in 500 mM KHCO3 (pH 8.3) and an Ar or CO2 atmosphere are shown in Fig. 4. In the presence of CO2, the current began to flow at a more positive potential (−1.2 V) than in Ar (−1.4 V). In this respect, these voltammograms are similar to those obtained by Kaneko et al. using a solid Au electrode in KOH+ methanol after purging the electrolyte with N2 or CO2 [26]. In the presence of Ar, CO was not detected but was detected after each voltammetric cycle measured in the presence of CO2. Whereas increasing the Ar pressure
Fig. 3. Schematic diagram of a pore in the electrode membrane. The region where the meniscus meets the Au catalyst is assumed to be the most active site for CO2 reduction. C, Pt counter electrode; W, Au working electrode; R, SCE reference electrode; M, porous polymer membrane; E, 500 mM KHCO3; S, current source; V, voltage source.
Fig. 5. The reaction order i of the rate-limiting step in CO2 reduction, calculated from the positive going (solid line) and negative going (dashed line) CO2 reduction current shown in the inset for CO2 reduction at 10 kPa (dashed line) and 50 kPa (solid line) above atmospheric pressure.
did not change the voltammogram, higher current densities were observed when the CO2 inlet pressure was increased from 10 to 50 kPa above atmospheric pressure. It is expected from the observations of Hara et al. [27], who obtained current densities of 163 mA cm − 2 at a solid Au electrode and a CO2 pressure of 30 atm in aqueous electrolyte, that significantly higher current densities could also be obtained at similarly high CO2 pressures. These observations indicate that CO formed as a result of the electrochemical reduction of CO2 rather than bicarbonate ion [7]. It follows that the difference between the current density in the presence of CO2 and Ar, the partial current density for CO2 reduction ( jCO2), is due to the reduction of dissolved CO2. This is proportional to the partial pressure of CO2 (pCO2) [9] jCO2 = kp iCO2
(4)
where the proportionality constant (k) and the reaction order of the rate-determining step (i), are functions of the potential. Solving Eq. (1) for i at two values of pCO2 gives i=
Fig. 4. Cyclic voltammograms of a porous Au electrode in Ar (dotted line) and CO2 at 10 kPa (dashed line) and 50 kPa (solid line) above atmospheric pressure, measured at a scan rate of 100 mV s − 1. The transient current density ( j ) was calculated with respect to the geometrical area of the electrode.
log(jCO2/ j %CO2) log(pCO2/p%CO2)
(5)
The reaction order for the rate-limiting step of CO2 reduction, calculated with Eq. (5) and the voltammograms in Fig. 4, is shown in Fig. 5. In the region between − 1.3 and −2.4 V, i is fractional, indicating that adsorbed CO2 takes part in the rate-limiting reaction in this potential range [14]. This supports the view that an adsorbed species such as CO2− (ads) is involved in CO2 reduction under these conditions. The onset potential for CO2 reduction at the thin film Au electrode, − 1.2 V, is comparable with that of an
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Au-loaded carbon GDE in 200 mM KHCO3 [23]. This onset potential for CO2 reduction is more negative than the reduction potential for CO2 to CO of − 0.84 V at pH 8.3 [3]. The high overpotential indicates initial electron transfer to form an intermediate product [10]. However, the onset potential is 890 or 970 mV more positive than the standard potentials for formation of the anion radical CO2− in aqueous media of − 2.09 V according to Surdhar et al. [28] or − 2.17 V according to Koppenal and Rush [29]. This may indicate that the adsorbed radical anion forms at a more positive potential than the fully solvated radical. Alternatively, as pointed out by Hirota et al. [30], who observed a similar effect using semiconductor electrodes in aprotic media, this can be explained by the formation of a radical anion complex, described by Amatore and Save´ ant [11] in a reaction scheme for CO2 reduction in solvents of low proton availability: 2CO2 (ads)+e − (CO2)2 − (ads)
(6)
On the basis of molecular beam experiments, and ignoring solvation effects, Hirota et al. calculated that the complex has a reduction potential of −0.64 V. If the radical complex does form, after further reduction by transfer of an electron, then it may dissociate into CO and CO23 − by disproportionation, according to Eq. (3). However, building on the scheme of Hori et al. [10], we suggest that a competing reaction in aqueous media is the reaction of H+ with an O atom from an adsorbed radical complex, forming ((CO2)2H) (ads). This may then be further reduced by another electron transfer to form CO2 (ads) with the release of CO and OH−: ((CO2)2H) (ads)+ e − CO2 (ads) + CO +OH −
(7)
Fig. 6. Composition of the effluent gas after passing CO2 over a series of eight porous Au electrodes at a range of applied potentials. The effluent gas at flow rates of between 225 and 350 ml s − 1 (filled) and 75 ml s − 1 (hollow) was analysed for CO , and CO2 , . The errors are the standard errors of the mean of five measurements. The residual gas, assumed to be H2, is also shown , .
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3.3. Composition of the effluent gas After passing CO2 over a series of eight porous Au electrodes, the main gas products detected using gas chromatography were CO and H2. In the FTIR spectra, absorption peaks from 2200 to 2100 cm − 1, which are characteristic of CO, confirm the presence of this gas in the effluent stream. The proportion of CO and CO2 detected in the effluent gas using gas chromatography and infra-red absorption and the calculated residual gas is shown in Fig. 6. Quantitative measurements of H2 could not be made with our equipment. However, H2 was the only other gas detected by GC, so the balance of the composition of the gas was assumed to be due to H2. Some negative values for the estimated rate of H2 formation might be due to slow release of CO2 from the walls of the sample holder that had adsorbed during earlier measurements when the CO2 concentration was relatively high. The results are divided into a group of outlet flow rates between 230 and 333 ml s − 1, which were analyzed by GC, and a second group at a more positive potential with an outlet flow rate of around 75 ml s − 1, which were analyzed by FTIR. In both groups, as the potential became more negative, the concentration of CO increased and that of CO2 decreased. At an applied potential of −2.2 V, the composition of the out-flowing gas was 21% CO and 77% CO2. The composition of this gas reaction product is significantly different from that obtained by Azuma et al. [13] who used a solid Au electrode and CO2 saturated 0.05 M KHCO3 at 0 °C. Under these conditions, the main reaction products obtained at − 2.2 V were 73.4% H2 and 16.9% CO. It has been suggested by Ohta et al. [31], that the current density during CO2 reduction at Cu electrodes affects the type and amount of reaction products. Unfortunately, Azuma et al. do not state the current density of the reaction, so it is not possible to know if this contributed to the difference between the relative amounts of our reaction products. At flow rates of between 225 and 350 ml s − 1, the composition of CO exceeded the CO2 concentration at − 3.4 V. As the potential became more negative the concentration of H2 increased. At the lower flow rate, the CO concentration exceeded the CO2 concentration at a more positive potential (− 1.7 V) than at the higher flow rate. Thus, by varying the flow rate and potential, it is possible to obtain a composition of outflowing gas that is required by the Fischer– Tropsch Process, which according to Klier et al. [32] has a maximum synthesis rate for CO2/CO/H2 ratios of 2/28/70.
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Fig. 7. Faradic efficiency (FCO) for CO2 reduction to CO at flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 . The data points have been qualitatively fitted with a curve. At potentials more negative than − 1.75 V, the Faradaic efficiencies at flow rates of 75 ml s − 1 are lower than those at higher flow rates because of a decrease in the partial pressure of CO2.
Fig. 8. Tafel plot of the partial current density for CO formation ( −jCO) at outlet gas flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 . Non-linear least squares fits to the first (dotted line), second (solid line) and third (dashed line) linear regions are also shown.
3.4. Faradaic efficiency The Faradaic efficiency for CO formation (FCO) was calculated assuming two electrons are required to form each CO molecule. As shown in Fig. 7, there is considerable scatter in the Faradaic efficiency for CO formation. Nevertheless, it increases from zero percent to over 67% in the range −1.2 to −2.2 V. The highest Faradaic efficiency for CO formation was 85%, which is comparable with that obtained by Hori et al. at solid Au electrodes [10]. These results are different from those obtained by Shibata et al. [23], using an Au loaded PTFE-bonded carbon GDE with 0.2 M KHCO3. In this case, the current efficiency for CO formation was 20% in the range −0.99 to − 1.34 V, after which it increased
sharply to 45%, maintaining this value until it reached − 2.5 V, decreasing as the potential became more negative. The current efficiency for HCOOH formation was fairly constant at about 10% over this range. In our case, the composition of the outflowing gas at − 2.0 V was mainly CO and CO2, with almost no H2. Therefore, it is expected that the Faradaic efficiency for HCOOH formation will be about 30–40% at this potential. Use of catalysts that further limit the formation of H2 and HCOOH may increase the Faradaic efficiency for CO formation. This may be achieved by selecting a Au crystal structure that favours CO formation [33], or by using additional catalysts for CO formation, such as Ni –cyclam complexes. These may be used either in solution or incorporated directly onto the surface of the electrode, as recently demonstrated by Jarzebinska et al. [34].
3.5. Partial current density for CO formation Fig. 8 shows a Tafel plot of the partial current density for CO formation ( jCO). At − 1.38 V, the partial current density is 5 mA cm − 2, which is comparable to that obtained at a massive Au electrode in 0.1 M KHCO3 (4.4 mA cm − 2) [10]. These plots are linear at applied potentials of between − 1.2 to − 1.5 V, − 1.7 to − 2.2 V and −2.7 to − 3.6 V. The general shape of the first two linear regions correspond to those obtained by Vassilev et al. during the electroreduction of CO2 at Sn, In, Pb and other metals in phosphate buffer [9]. Here, it was claimed that the rate-determining step in the first linear (Tafel) region is the transfer of an electron to adsorbed CO2 radical anions to form CO. They ascribe the second linear region to the initial formation of an adsorbed CO2 radical anion by transfer of an electron to an adsorbed CO2 molecule. The very high slope of this linear region is possibly symptomatic of the expected diffusion, rather than activation, controlled kinetics for this reaction. Although the potentials have not been corrected for ohmic losses, the slope of the first region (556 mV per decade increase in current density) is comparable to that obtained at a Pb-impregnated PTFE-bonded carbon GDE (500 mV decade − 1) [17]. In this case, the high Tafel slope was thought to be due to the formation of a non-conducting barrier. Such a barrier may be formed from adsorbed CO− 2 radical anions. The formation of such a barrier would account for the fractional order of the reaction with respect of CO2 pressure [9] in this potential range, as is shown in Fig. 5. The linearity of the second region, with a slope of 1.87 V decade − 1 increase in current, probably indicates control by either gas diffusion or ionic mass transport, whereas the slope of the third region (9.79 V decade − 1) is probably due to the combined effects of the two [21].
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The capacity of the electrode to store electrical power by converting CO2 into CO (SCO) is proportional to the partial current density for CO formation: SCO =
jCODH uCO nF
(8)
The calculated values of Eq. (8), shown in Fig. 9, demonstrate that as the potential becomes more negative, the power storage capacity steadily increases from about 50 W cm − 2 at − 1.2 V to over 300 W m − 2 at − 2.5 V.
3.6. Energy con6ersion efficiency The first law on enthalpic efficiency [35] for converting electrical energy into the enthalpy of CO was calculated assuming the reaction scheme in Fig. 3, in which there are no changes in the chemical potential of
Fig. 9. The capacity of the electrode for storage of electrical power by converting CO2 into CO (SCO), calculated using Eq. (8), at outlet flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 .
Fig. 10. First law efficiency of cell for converting electrical energy to enthalpy of CO (mCO) at flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 , calculated using Eq. (9). The energy efficiency, calculated using Eq. (10) and the line of best fit in Fig. 7, is shown for comparison (solid line).
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the electrolyte. In this case, the first law energy conversion efficiency (mCO) is equivalent to the ratio of the rate at which enthalpy is stored in CO to the rate of electrical energy used to produce it: mCO =
fCODH uCO (EW − EC)
(9)
where fCO is the rate of formation of the CO and DH uCO is the standard molar enthalpy change when CO is burnt in oxygen and EC is the potential of the counter electrode. The first law efficiency, calculated from Eq. (9) is shown in Fig. 10. An alternative expression for the overall energy conversion efficiency of the cell is the product of the Faradaic efficiency and the ideal enthalpic efficiency: mCO =
FCODH uCO nF(EW − EC)
(10)
where n is the number of electrons required in the electrochemical reduction of CO2 and F is Faraday’s constant. Eq. (10) is the inverse of the energy efficiency of a fuel cell [36]. As is shown in Fig. 10, the energy efficiency as calculated from Eq. (10) is in good agreement with that obtained using Eq. (9). At −1.5 V the energy storing efficiency was greater than 25% when calculated by both methods. According to the trend in Fig. 10, the maximum efficiency will be obtained as the applied potential approaches the onset potential. Eqs. (9) and (10) show that the energy storage efficiency can be increased by increasing either the Faradaic efficiency for CO formation or by decreasing the potential difference between the working and counter electrodes. Part of the energy lost during the formation of CO was observed as an increase in electrolyte temperature. Other losses include H2 and formate production at the cathode and hydroxyl and formate oxidation at the anode [37]. We measured the impedance spectrum of the eight cell unit, to estimate the energy lost due to electrode and electrolyte resistance. Before taking the measurements, the electrolyte was saturated with CO2 by passing the gas through the cells for several hours. The impedance spectra are shown in Fig. 11, along with the calculated frequency dispersion of an equivalent circuit. Using the notation of Boukamp [38], the equivalent circuit which gave a good fit to the measured values was a resistor (R1), a capacitor (C), an inductor (L) and a constant phase element (CPE: Q1, n1), all in series with a resistor (R2) and CPE (Q2, n2) in parallel. The values of the parameters of the equivalent circuit which gave a good fit to the impedance spectrum were R1 = 1.89 V, C=8.7× 10 − 4 F, Q1 = 2.1× 10 − 3, n1 = 0.81, L=1.6× 10 − 6 H, R2 = 0.35 V, Q2 = 6.6×10 − 3 and n2 = 0.86. From this analysis, the combined electrode and electrolyte resistance was 1.89 V. During CO2 reduction at − 1.5 V versus SCE, the potential difference between the working and counter electrode was
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Fig. 11. The magnitude ( z ) and the phase angle (h) of an impedance spectrum of the eight cell unit, obtained at an applied voltage of 5 mV rms. The electrolyte was 500 mM KHCO3 and the reference electrode was the counter electrode during this measurement. Also shown is z (solid) and h (dashed) from the calculated frequency dispersion of an equivalent circuit.
Fig. 12. Time dependence of current density during electroreduction of CO2 at a Au film cathode using a Pt anode in 500 mM KHCO3. At an applied potential of − 2.1 V, the gas flow rate was 8 ml s − 1 . At an applied potential of − 1.7 V, the gas flow rate was 33 ml s − 1
.
Under these conditions, the stability of this electrode was comparable to Dewulf and Bard’s SPE-supported GDE [19]. The resistance of the Au film remained unchanged during the experiment. The decline in current density may be due to one or a combination of factors, including poisoning of the electrode from impurities in the electrolyte or feed gas or from the reaction products. It is possible that CO may have been reduced to elemental carbon, lowering the number of available CO2 binding sites and resulting in a decrease in the current density. Alternatively, a trace metal in the KHCO3 such as Fe could have plated onto the Au electrode. If this occurred, the percentage of H2 produced would increase at the expense of CO [13]. Since adsorbed CO is expected to chemisorb at sites normally occupied by H2 [7], a decline in the current density could result from a decrease in H2 formation that would otherwise have occurred at sites occupied by CO. At an applied potential of − 1.7 V and a gas flow rate of 33 ml s − 1, the current density stabilized at around 8 mA cm − 2, remaining at this level for 48 h. After this time, the main reaction product was a gas other than CO. Based on the colour of the flame when this gas was burned in air, it was assumed to be H2. The formation of H2 supports the suggestion that metals like Fe may have plated onto the Au film. The decrease in current density over several days was accompanied by the formation of a blackish film, which was more pronounced near the exit of the electrode than at the entrance. This suggests the formation of a carbon deposit due to the reduction of CO. The formation of a carbon deposit on the electrode would probably also contribute to the observed decrease in the current density.
4. Conclusions 3.85 V and the total current flowing through the eight cell unit was −340 mA. Under these conditions, a potential drop of 646 mV between the working and counter electrodes was due to the resistance of the electrolyte and electrode, therefore a 5% improvement in the energy efficiency would be expected if the resistance of the electrode and electrolyte was eliminated.
3.7. Electrode stability In order to obtain an indication of the stability of the electrode over time, the current was measured as CO2 was passed over the series of eight Au electrodes. As shown in Fig. 12, at −2.1 V and a gas flow rate of 8 ml s − 1, the current rose to above 25 mA cm − 2. However, after 10 min the current density decreased by 40% over the following 100 min.
This work demonstrates the use of a porous thin Au film electrode for the electrochemical reduction of CO2 to CO, in aqueous KHCO3. The hydrophilic property of the supporting polymer substrate for the Au film was crucial for obtaining current densities comparable with those obtained at ordinary Au electrodes. Cyclic voltammetry in the presence and absence of CO2 confirmed that dissolved CO2 was reduced, rather than bicarbonate. The reaction kinetics were of fractional order in CO2 partial pressure, indicating that the rate-limiting step of CO2 reduction involves a bound species, possibly CO2−. However, the onset potential for CO2 reduction was about 900 mV more positive than the standard reduction potential for CO2− formation in aqueous media. Possible reasons for this were suggested.
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The Faradaic efficiency for CO2 reduction to CO was over 75%, which is comparable to that obtained at an Au-loaded GDE. The energy efficiency for energy storage in CO rose to a maximum of 35% near the onset potential for CO2 reduction. Impedance spectra of the electrodes indicated that up to a 5% increase in the energy efficiency could be expected if the electrode and electrolyte resistance was eliminated. Apart from heating and less than optimal Faradaic efficiency, the remaining energy losses are assumed to be due to the oxidation of COOH− and OH−. The current density during the electrochemical reduction of CO2 was stable for 48 h at an applied potential of − 1.7 V. At − 2.3 V, a steady decline in the current density was observed. Various reasons for the decrease in the current density were suggested.
Acknowledgements We thank Drs J.J. Lowke, D.L. Tilbrook and V. Braach-Maksvytis for helpful discussions and Drs R. Driver, C.P. Foley and A.F. Collings for critical reading of the manuscript. We also thank Drs I.C. Plumb, L.M. Besley and Mr M. Arnautovitch for assistance with gas analysis, Dr A. Molodyk for analysis of XRD spectra, Dr S. Lam for the SEM image and Mr G.R. Baxter for help in making the metal films.
References [1] M.M. Halmann, Chemical Fixation of Carbon Dioxide: Methods for Recycling CO2 into Useful Products, CRC Press, Boca Raton, FL, 1993 Ch. 1. [2] M.A. Scibioh, V.R. Vijayaraghavan, J. Sci. Ind. Res. 57 (1998) 111. [3] F.R. Keene, in: B.P. Sullivan, K. Krist, H.E. Guard (Eds.), Electrochemical and Electrocatalytic Reactions of Carbon Dioxide, Elsevier, Amsterdam, 1993 Ch. 1. [4] J.J. Lowke, A.F. Collings, V. Braach-Maksvytis, R. Pace, J. Andrews, B. Osmond, in: Proceedings of the Fifth International Conference on Greenhouse Gas Control Technologies, CSIRO, 2000, p. 624. [5] M. Aresta, in: M. Aresta, D. Forti (Eds.), Carbon Dioxide as a Source of Carbon, Reidel Publishing Co, Dordrecht, 1987 Ch. 1. [6] C. O’Connell, S.I. Hommeltoft, R. Eisenberg, in: M. Aresta, D. Forti (Eds.), Carbon Dioxide as a Source of Carbon, Reidel Publishing Co, Dordrecht, 1987 Ch. 3.
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[7] K.W. Freese Jr., in: B.P. Sullivan, K. Krist, H.E. Guard (Eds.), Electrochemical and Electrocatalytic Reactions of Carbon Dioxide, Elsevier, Amsterdam, 1993 Ch. 6. [8] A.W.B. Aylmer-Kelly, A. Bewick, P.R. Cantrill, A.M. Tuxford, Faraday Discuss. Chem. Soc. 56 (1974) 96. [9] Yu.B. Vassiliev, V.S. Bagotzky, N.V. Osetrova, O.A. Khazova, N.A. Mayorova, J. Electroanal. Chem. 189 (1985) 271. [10] Y. Hori, H. Wakebe, T. Tsukamoto, O. Koga, Electrochim. Acta 39 (1994) 1833. [11] C. Amatore, J.-M. Save´ ant, J. Am. Chem. Soc. 103 (1991) 5021. [12] D.J. Schiffrin, Faraday Discuss. Chem. Soc. 56 (1974) 75. [13] M. Azuma, K. Hashimoto, M. Hiramoto, M. Watanabe, T. Sakata, J. Electrochem. Soc. 137 (1990) 1772. [14] Yu.B. Vassiliev, V.S. Bagotzky, O.A. Khazova, N.A. Mayorava, J. Electroanal. Chem. 189 (1985) 295. [15] D.P. Summers, S. Leach, K.W. Frese Jr., J. Electroanal. Chem. 205 (1986) 219. [16] R.L. Cook, R.C. MacDuff, A.F. Sammells, J. Electrochem. Soc. 137 (1990) 607. [17] M.N. Mahmood, C. Masheder, C.J. Harty, J. Appl. Electrochem. 17 (1987) 1159. [18] E. Mendez, M.E. Martins, C.F. Zinola, J. Electroanal. Chem. 477 (1999) 41. [19] D.W. Dewulf, A.J. Bard, Catal. Lett. 1 (1988) 73. [20] R.L. Cook, R.C. MacDuff, A.F. Sammells, J. Electrochem. Soc. 135 (1988) 1470. [21] M.L. Perry, J. Newman, E.J. Cairns, J. Electrochem. Soc. 145 (1998) 5. [22] N. Chatterjee, I.N. Basumallick, J. New Mat. Electrochem. Sys. 2 (1999) 157. [23] M. Shibata, K. Yoshida, N. Furuya, J. Electroanal. Chem. 442 (1998) 67. [24] J.O.’M. Bockris, B.D. Cahan, J. Chem. Phys. 50 (1969) 1307. [25] Millipore Product Information. [26] S. Kaneko, K. Iiba, K. Ohta, T. Mizuno, A. Saji, J. Electroanal. Chem. 441 (1998) 215. [27] K. Hara, A. Kudo, T. Sakata, J. Electroanal. Chem. 391 (1995) 141. [28] P.S. Surdhar, S.P. Mezyk, D.A. Armstrong, J. Phys. Chem. 93 (1989) 3360. [29] W.H. Koppenol, J.D. Rush, J. Phys. Chem. 91 (1987) 4429. [30] K. Hirota, D.A. Tryk, T. Yamamoto, K. Hashimoto, M. Okawa, A. Fujishima, J. Phys. Chem. B 102 (1998) 9834. [31] K. Ohta, A. Hasimoto, T. Mizuno, Energy Convers. Management 36 (1995) 625. [32] K. Klier, V. Chatikavanij, R.G. Herman, G.W. Simmons, J. Catal. 74 (1982) 343. [33] N. Hoshi, M. Noma, T. Suzuki, Y. Hori, J. Electroanal. Chem. 421 (1997) 15. [34] A. Jarzebinska, P. Rowinski, I. Zawisza, R. Bileweicz, L. Siegfried, T. Kaden, Anal. Chim. Acta 396 (1999) 1. [35] M.G. Patterson, Energy Policy 24 (1996) 377. [36] T. Erdey-Gru´ z, Kinetics of Electrode Processes, Adam Hilger Ltd, London, 1973. [37] R. Williams, R.S. Crandell, A. Bloom, Appl. Phys. Lett. 33 (1978) 381. [38] B.A. Boukamp, Equivalent Circuit Users Manual, University of Twente, 1989.
Energy storage by the electrochemical reduction of CO2 to CO at a porous Au film Gregory B. Stevens *, Torsten Reda, Burkhard Raguse CSIRO Di6ision of Telecommunications and Industrial Physics, PO Box 218, Lindfield, NSW 2070, Australia Received 16 August 2001; received in revised form 4 January 2002; accepted 6 January 2002
Abstract We have investigated the performance of a porous Au film electrode for future applications in energy storage by the reduction of CO2 to CO. The electrode consists of a porous hydrophilic polymer membrane with a 260 nm Au film, applied by vapour deposition. The hydrophilic property of the polymer substrate was necessary for obtaining current densities comparable to those at an ordinary Au electrode under similar conditions: aqueous KHCO3, ambient temperature and approximately one atmosphere pressure of CO2. At potentials more negative than the activation potential for CO2 reduction, − 1.2 V versus SCE, the current density was of fractional order with respect to the CO2 pressure, indicating the involvement of an adsorbed reaction intermediate, possibly a radical anion complex (CO2)2−. Analysis by GC and FTIR of the gas reaction product showed that the energy storage efficiency was maximal near the activation potential. Close to this potential, CO was formed with a Faradaic efficiency of 75% and an enthalpic efficiency of 35%. Under these conditions, the power storage capacity of the electrode was 50 W m − 2. © 2002 Elsevier Science B.V. All rights reserved. Keywords: Carbon dioxide; Au electrode; Carbon monoxide; Energy; Electrocatalytic reduction; Current efficiency
1. Introduction As natural oil reserves become depleted, attention will increasingly focus on alternative energy sources. These may include hydrocarbons and oxygenated aliphatics, synthesized from CO and H2 via the Fischer –Tropsch process. The energy and raw materials for this are currently derived from the burning of coal, with the accompanying release of CO2 as a by-product. However, concerns that increases in atmospheric CO2 may lead to serious global climate changes have led a number of authors to suggest that CO2 itself may be used as a source of carbon for the production of petroleum-like materials [1 – 3]. This may then lead to the possibility of regulating the concentration of atmospheric CO2 [4,5]. As CO2 is one of the most thermodynamically stable carbon compounds, a highly energetic reductant or an external source of energy is required to convert it into * Corresponding author. Tel.: +61-29413-7490; fax: + 61-294137202. E-mail address: [email protected] (G.B. Stevens).
other carbon compounds. It is well known that CO2 can be reduced electrochemically [6,7]. The mechanism of this begins with the formation of an adsorbed CO2 radical anion intermediate [8]: CO2 (ads)+e − CO2 − (ads)
(1)
The dominant reaction, subsequent to the initial formation of CO2− (ads), depends on factors such as the type and pH of the electrolyte, the electrode potential, stirring, the CO2 partial pressure and the electrode material [9]. According to Hori et al. [10], in aqueous media transfer of a proton from water to an adsorbed radical anion results in the formation of CO: CO2 − (ads)+H+ + e − CO + OH −
(2)
This is catalysed by metals in the order Au\ Ag \Zn. In aprotic electrolytes, adsorbed radical anions combine with CO2, resulting in the formation of CO and CO23 − by disproportionation [11]: 2CO2 − + e − CO + CO23 −
0022-0728/02/$ - see front matter © 2002 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 2 - 0 7 2 8 ( 0 2 ) 0 0 6 8 8 - 5
(3)
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At other metals from Groups IIB, IIIB and IVB in aqueous electrolytes, formate is the dominant reaction product in the order of activity Cd\ Sn\ In \ Pb \ Tl \Hg [12,13]. This is due to weak binding of CO2 radical anions at these metals. After forming at the electrode, free radical anions are protonated to produce formate ions, whereas in aprotic solvents, free radical anions dimerize to form oxalate [14,10]. In acids, CO2 is reduced to methanol on Mo [15] and to hydrocarbons at Cu electrodes [16]. Other metals show little activity in reducing CO2, either because the hydrogen overpotential is too low or because CO binds tightly to the metal, preventing further reactions with CO2 from taking place. The low solubility of CO2 in aqueous electrolytes is a limiting factor in the reaction rate and, as a result, the reaction will be diffusion limited. The mass transport limitation can be overcome to some extent by the use of gas diffusion electrodes (GDEs), which interface the gas, electrocatalyst and electrolyte phases [1]. Typically, GDEs are constructed from metal-impregnated polytetrafluoroethylene (PTFE)-bonded carbon powders [17], in which the CO2 diffuses through a porous hydrophobic region in the GDE until it reaches the electrolyte in a hydrophilic region, where it is reduced [18]. Although high current densities can be obtained with this type of electrode, the hydrophobic region is prone to flooding. Cook et al. report that this results in a gradual decline of the current density [16]. This problem may be overcome by the use of a solid polymer electrolyte (SPE) such as a perfluorosulfonate ion-exchange resin (Nafion®). Such electrodes have been used in the reduction of CO2 to CH4 and C2H4 [19,20]. However, Nafion®-based electrodes work best in acid media and only relatively small Faradaic efficiencies and current densities have been achieved in the reduction of CO2 with these GDEs. To minimize the adverse effects of mass transport in GDEs, the thickness of the active layer should be kept to a minimum [21] which also leads to economic utiliza-
tion of expensive electrocatalysts. Recently, Chatterjee and Basumalick reported the use of a lightweight GDE made from a thin Ag film supported on a porous PVC membrane, for use in fuel cells [22]. Because this type of GDE can be constructed with an inert substrate, it can be used with electrolytes in which CO2 is more soluble. Although Au is highly selective for CO formation, it has not been studied in great depth for this purpose. Shibata et al. investigated the use of an Au impregnated PTFE-bonded carbon GDE for the electrochemical reduction of CO2 and nitrate to urea [23]. In this work we investigate the electrochemical reduction of CO2 to CO in aqueous media at a thin porous Au GDE and consider the potential application of this system for energy storage.
2. Experimental Throughout this paper, the potential of the working electrode (EW) is measured with respect to a standard calomel electrode (SCE) and current densities are expressed in terms of the geometrical area of the electrode.
2.1. Electrode preparation The film electrodes were prepared by vapour deposition of Au onto one side of porous membranes made of polyvinylidene fluoride polymer (Millipore GVWP), polyethylene (DSM 7P20) and PTFE (Millipore FGLP). Gold (99.99% Au) was evaporated under vacuum (Edwards model 12EA/461), directly onto membranes located 30 cm above an electrically heated tungsten boat. The evaporation rate was 0.2–1.0 nm s − 1 and was continued until the final thickness was 200– 260 nm, as measured with a thickness monitor (Maxtek TM-100). The crystal structure of the metal film was investigated with an X-ray diffraction spectrometer (Philips X’pert). Microscopy and microanalysis was done with a scanning electron microscope (Jeol JSM-6300F) equipped with an energy-dispersive X-ray spectrometer.
2.2. Cell and electrolyte
Fig. 1. Schematic diagram of the electrochemical cell. R, saturated calomel reference electrode; C, Pt wire coil counter electrode; E, 500 mM KHCO3 electrolyte; W, working electrode; G, gold leaf contact to the working electrode; O, ring seal; P, inlet pressure gauge; F, outlet pressure gauge and flow meter.
The film electrodes were supported in an electrolysis cell illustrated in Fig. 1. The cell body was made of borosilicate glass and the gas cavity of methyl methacrylate polymer. Electrical connection to the Au film was made using Au foil. A nitrile ‘O’ ring between the membrane and the cell housing formed a fluid seal, leaving an electrode area of 3.17 cm − 2. Cyclic voltammograms were measured with an electrochemistry recording unit (AD Instruments).
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Gas samples were injected into a Porapak Q column in a gas chromatograph equipped with a thermal conductivity detector (3400 GC Varian Inc.). Gas products were identified by comparison of the retention times with those of a reference gas. Quantitative measurements of CO and CO2 in the gas mixtures were obtained as the ratio of the peak areas of gas reaction products and the same volume of pure gas samples. The effluent gas was also analysed by measuring the absorption in the infra-red region using a Fourier transform infra-red spectrometer (BioRad FTS-65). The absorption integrated from 2200 to 2100 cm − 1 and from 2277 to 2255 cm − 1 was used to measure CO and CO2 concentrations in the gas outflow.
Fig. 2. Scanning tunnelling microscope image of the surface of the porous Au layer. The lighter regions indicate Au. EDX microanalysis did not show the presence of O or C, indicating that Au completely covered the substrate layer.
Measurements of the rate of CO formation were made after passing CO2 over a series of eight Au/ GVWP cathodes, so as to increase the volume of reaction products for analysis by gas chromatography and infra-red absorption. In this case, the membranes were supported in a methyl methacrylate polymer housing. A common counter electrode made of a Pt wire (0.5 mm dia× 540 mm) was located 5 mm from the membranes. Impedance spectra of the eight cell unit were measured using an electrochemical impedance system (EG&G Instruments model 263). The electrolyte was made from 99.99% KHCO3 and MilliQ water. The electrolyte was prevented from flowing into the gas phase by maintaining the gas pressure at 10 kPa above atmospheric pressure. Gasses used were Ar (99.99%) and CO2 (99.95%), which were supplied at pressures of between 10 and 50 kPa and at flow rates of up to 20 ml min − 1 via copper and FEP plastic tubing.
2.3. Electrolysis and analysis of products In the cyclic voltammetry work, traces of CO in the effluent gas were detected with a colorimetric CO detector (Onix Process Analysis). For the quantitative analysis of gas products, a current was passed for periods of up to 3 min at potentials of between − 1.0 and − 3.6 V. The electrolyte temperature was maintained below 26 °C by allowing for cooling between measurements. Gas products from these cells were passed through a 9 cm column of powdered silica gel to exclude water vapour before analysis. The gas flow rates were calculated from the time taken to collect a volume of gas at atmospheric pressure.
3. Results and discussion
3.1. Electrode structure The electrical resistance of the films was measured by positioning two Au bars, each 1 cm long and 2 mm wide, on opposite sides of an imaginary 1 cm square on the film surface. The resistance between the bars on these surfaces was between 1.6 and 2.0 V. However, when used with the KHCO3 electrolyte, Au electrodes made on hydrophobic membranes (DSM 7P20 and FGLP) had much lower current densities than those made on hydrophilic membranes (GVWP). The relatively high current density obtained with the GVWP membranes was attributed to its hydrophilicity. Fig. 2 shows an SEM image of the surface of the porous Au film evaporated onto a GVWP membrane. The surface appears to be sponge-like with pores of varying sizes of up to 1 mm across. The larger pores are much larger than the nominal pore diameter of the substrate layer, 220 nm. This suggests that structural changes to the GVWP membrane occurred during the evaporation, resulting in the Au film bonding well to the GVWP membrane. This was evident when sticky tape was applied to the Au layer and did not remove it when peeled off. The X-ray diffraction spectra of the film indicated the presence of polycrystalline Au, with a preferred (111) orientation and an average crystallite size of about 60 nm. Fig. 3 shows a schematic of the cross section of a membrane pore to indicate the assumed geometry of the gas electrode electrolyte interface. According to Bockris and Cahan’s Finite Contact Angle Meniscus model of porous electrodes [24], most of the current associated with the reduction of the gas phase will be produced very close to the meniscus tip. The electrolyte contacted the gas phase after passing through the membrane by capillary action. As the membrane became wet, it was assumed that the total surface energy of the water and membrane decreased
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and would need to be raised before gas will pass through to the electrolyte. This will occur at pressures above 340 kPa for the GVWP membrane [25].
3.2. Transient current-6oltage characteristics Cyclic voltammograms of the Au electrode in 500 mM KHCO3 (pH 8.3) and an Ar or CO2 atmosphere are shown in Fig. 4. In the presence of CO2, the current began to flow at a more positive potential (−1.2 V) than in Ar (−1.4 V). In this respect, these voltammograms are similar to those obtained by Kaneko et al. using a solid Au electrode in KOH+ methanol after purging the electrolyte with N2 or CO2 [26]. In the presence of Ar, CO was not detected but was detected after each voltammetric cycle measured in the presence of CO2. Whereas increasing the Ar pressure
Fig. 3. Schematic diagram of a pore in the electrode membrane. The region where the meniscus meets the Au catalyst is assumed to be the most active site for CO2 reduction. C, Pt counter electrode; W, Au working electrode; R, SCE reference electrode; M, porous polymer membrane; E, 500 mM KHCO3; S, current source; V, voltage source.
Fig. 5. The reaction order i of the rate-limiting step in CO2 reduction, calculated from the positive going (solid line) and negative going (dashed line) CO2 reduction current shown in the inset for CO2 reduction at 10 kPa (dashed line) and 50 kPa (solid line) above atmospheric pressure.
did not change the voltammogram, higher current densities were observed when the CO2 inlet pressure was increased from 10 to 50 kPa above atmospheric pressure. It is expected from the observations of Hara et al. [27], who obtained current densities of 163 mA cm − 2 at a solid Au electrode and a CO2 pressure of 30 atm in aqueous electrolyte, that significantly higher current densities could also be obtained at similarly high CO2 pressures. These observations indicate that CO formed as a result of the electrochemical reduction of CO2 rather than bicarbonate ion [7]. It follows that the difference between the current density in the presence of CO2 and Ar, the partial current density for CO2 reduction ( jCO2), is due to the reduction of dissolved CO2. This is proportional to the partial pressure of CO2 (pCO2) [9] jCO2 = kp iCO2
(4)
where the proportionality constant (k) and the reaction order of the rate-determining step (i), are functions of the potential. Solving Eq. (1) for i at two values of pCO2 gives i=
Fig. 4. Cyclic voltammograms of a porous Au electrode in Ar (dotted line) and CO2 at 10 kPa (dashed line) and 50 kPa (solid line) above atmospheric pressure, measured at a scan rate of 100 mV s − 1. The transient current density ( j ) was calculated with respect to the geometrical area of the electrode.
log(jCO2/ j %CO2) log(pCO2/p%CO2)
(5)
The reaction order for the rate-limiting step of CO2 reduction, calculated with Eq. (5) and the voltammograms in Fig. 4, is shown in Fig. 5. In the region between − 1.3 and −2.4 V, i is fractional, indicating that adsorbed CO2 takes part in the rate-limiting reaction in this potential range [14]. This supports the view that an adsorbed species such as CO2− (ads) is involved in CO2 reduction under these conditions. The onset potential for CO2 reduction at the thin film Au electrode, − 1.2 V, is comparable with that of an
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Au-loaded carbon GDE in 200 mM KHCO3 [23]. This onset potential for CO2 reduction is more negative than the reduction potential for CO2 to CO of − 0.84 V at pH 8.3 [3]. The high overpotential indicates initial electron transfer to form an intermediate product [10]. However, the onset potential is 890 or 970 mV more positive than the standard potentials for formation of the anion radical CO2− in aqueous media of − 2.09 V according to Surdhar et al. [28] or − 2.17 V according to Koppenal and Rush [29]. This may indicate that the adsorbed radical anion forms at a more positive potential than the fully solvated radical. Alternatively, as pointed out by Hirota et al. [30], who observed a similar effect using semiconductor electrodes in aprotic media, this can be explained by the formation of a radical anion complex, described by Amatore and Save´ ant [11] in a reaction scheme for CO2 reduction in solvents of low proton availability: 2CO2 (ads)+e − (CO2)2 − (ads)
(6)
On the basis of molecular beam experiments, and ignoring solvation effects, Hirota et al. calculated that the complex has a reduction potential of −0.64 V. If the radical complex does form, after further reduction by transfer of an electron, then it may dissociate into CO and CO23 − by disproportionation, according to Eq. (3). However, building on the scheme of Hori et al. [10], we suggest that a competing reaction in aqueous media is the reaction of H+ with an O atom from an adsorbed radical complex, forming ((CO2)2H) (ads). This may then be further reduced by another electron transfer to form CO2 (ads) with the release of CO and OH−: ((CO2)2H) (ads)+ e − CO2 (ads) + CO +OH −
(7)
Fig. 6. Composition of the effluent gas after passing CO2 over a series of eight porous Au electrodes at a range of applied potentials. The effluent gas at flow rates of between 225 and 350 ml s − 1 (filled) and 75 ml s − 1 (hollow) was analysed for CO , and CO2 , . The errors are the standard errors of the mean of five measurements. The residual gas, assumed to be H2, is also shown , .
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3.3. Composition of the effluent gas After passing CO2 over a series of eight porous Au electrodes, the main gas products detected using gas chromatography were CO and H2. In the FTIR spectra, absorption peaks from 2200 to 2100 cm − 1, which are characteristic of CO, confirm the presence of this gas in the effluent stream. The proportion of CO and CO2 detected in the effluent gas using gas chromatography and infra-red absorption and the calculated residual gas is shown in Fig. 6. Quantitative measurements of H2 could not be made with our equipment. However, H2 was the only other gas detected by GC, so the balance of the composition of the gas was assumed to be due to H2. Some negative values for the estimated rate of H2 formation might be due to slow release of CO2 from the walls of the sample holder that had adsorbed during earlier measurements when the CO2 concentration was relatively high. The results are divided into a group of outlet flow rates between 230 and 333 ml s − 1, which were analyzed by GC, and a second group at a more positive potential with an outlet flow rate of around 75 ml s − 1, which were analyzed by FTIR. In both groups, as the potential became more negative, the concentration of CO increased and that of CO2 decreased. At an applied potential of −2.2 V, the composition of the out-flowing gas was 21% CO and 77% CO2. The composition of this gas reaction product is significantly different from that obtained by Azuma et al. [13] who used a solid Au electrode and CO2 saturated 0.05 M KHCO3 at 0 °C. Under these conditions, the main reaction products obtained at − 2.2 V were 73.4% H2 and 16.9% CO. It has been suggested by Ohta et al. [31], that the current density during CO2 reduction at Cu electrodes affects the type and amount of reaction products. Unfortunately, Azuma et al. do not state the current density of the reaction, so it is not possible to know if this contributed to the difference between the relative amounts of our reaction products. At flow rates of between 225 and 350 ml s − 1, the composition of CO exceeded the CO2 concentration at − 3.4 V. As the potential became more negative the concentration of H2 increased. At the lower flow rate, the CO concentration exceeded the CO2 concentration at a more positive potential (− 1.7 V) than at the higher flow rate. Thus, by varying the flow rate and potential, it is possible to obtain a composition of outflowing gas that is required by the Fischer– Tropsch Process, which according to Klier et al. [32] has a maximum synthesis rate for CO2/CO/H2 ratios of 2/28/70.
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Fig. 7. Faradic efficiency (FCO) for CO2 reduction to CO at flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 . The data points have been qualitatively fitted with a curve. At potentials more negative than − 1.75 V, the Faradaic efficiencies at flow rates of 75 ml s − 1 are lower than those at higher flow rates because of a decrease in the partial pressure of CO2.
Fig. 8. Tafel plot of the partial current density for CO formation ( −jCO) at outlet gas flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 . Non-linear least squares fits to the first (dotted line), second (solid line) and third (dashed line) linear regions are also shown.
3.4. Faradaic efficiency The Faradaic efficiency for CO formation (FCO) was calculated assuming two electrons are required to form each CO molecule. As shown in Fig. 7, there is considerable scatter in the Faradaic efficiency for CO formation. Nevertheless, it increases from zero percent to over 67% in the range −1.2 to −2.2 V. The highest Faradaic efficiency for CO formation was 85%, which is comparable with that obtained by Hori et al. at solid Au electrodes [10]. These results are different from those obtained by Shibata et al. [23], using an Au loaded PTFE-bonded carbon GDE with 0.2 M KHCO3. In this case, the current efficiency for CO formation was 20% in the range −0.99 to − 1.34 V, after which it increased
sharply to 45%, maintaining this value until it reached − 2.5 V, decreasing as the potential became more negative. The current efficiency for HCOOH formation was fairly constant at about 10% over this range. In our case, the composition of the outflowing gas at − 2.0 V was mainly CO and CO2, with almost no H2. Therefore, it is expected that the Faradaic efficiency for HCOOH formation will be about 30–40% at this potential. Use of catalysts that further limit the formation of H2 and HCOOH may increase the Faradaic efficiency for CO formation. This may be achieved by selecting a Au crystal structure that favours CO formation [33], or by using additional catalysts for CO formation, such as Ni –cyclam complexes. These may be used either in solution or incorporated directly onto the surface of the electrode, as recently demonstrated by Jarzebinska et al. [34].
3.5. Partial current density for CO formation Fig. 8 shows a Tafel plot of the partial current density for CO formation ( jCO). At − 1.38 V, the partial current density is 5 mA cm − 2, which is comparable to that obtained at a massive Au electrode in 0.1 M KHCO3 (4.4 mA cm − 2) [10]. These plots are linear at applied potentials of between − 1.2 to − 1.5 V, − 1.7 to − 2.2 V and −2.7 to − 3.6 V. The general shape of the first two linear regions correspond to those obtained by Vassilev et al. during the electroreduction of CO2 at Sn, In, Pb and other metals in phosphate buffer [9]. Here, it was claimed that the rate-determining step in the first linear (Tafel) region is the transfer of an electron to adsorbed CO2 radical anions to form CO. They ascribe the second linear region to the initial formation of an adsorbed CO2 radical anion by transfer of an electron to an adsorbed CO2 molecule. The very high slope of this linear region is possibly symptomatic of the expected diffusion, rather than activation, controlled kinetics for this reaction. Although the potentials have not been corrected for ohmic losses, the slope of the first region (556 mV per decade increase in current density) is comparable to that obtained at a Pb-impregnated PTFE-bonded carbon GDE (500 mV decade − 1) [17]. In this case, the high Tafel slope was thought to be due to the formation of a non-conducting barrier. Such a barrier may be formed from adsorbed CO− 2 radical anions. The formation of such a barrier would account for the fractional order of the reaction with respect of CO2 pressure [9] in this potential range, as is shown in Fig. 5. The linearity of the second region, with a slope of 1.87 V decade − 1 increase in current, probably indicates control by either gas diffusion or ionic mass transport, whereas the slope of the third region (9.79 V decade − 1) is probably due to the combined effects of the two [21].
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The capacity of the electrode to store electrical power by converting CO2 into CO (SCO) is proportional to the partial current density for CO formation: SCO =
jCODH uCO nF
(8)
The calculated values of Eq. (8), shown in Fig. 9, demonstrate that as the potential becomes more negative, the power storage capacity steadily increases from about 50 W cm − 2 at − 1.2 V to over 300 W m − 2 at − 2.5 V.
3.6. Energy con6ersion efficiency The first law on enthalpic efficiency [35] for converting electrical energy into the enthalpy of CO was calculated assuming the reaction scheme in Fig. 3, in which there are no changes in the chemical potential of
Fig. 9. The capacity of the electrode for storage of electrical power by converting CO2 into CO (SCO), calculated using Eq. (8), at outlet flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 .
Fig. 10. First law efficiency of cell for converting electrical energy to enthalpy of CO (mCO) at flow rates of between 225 and 350 ml s − 1 , and 75 ml s − 1 , calculated using Eq. (9). The energy efficiency, calculated using Eq. (10) and the line of best fit in Fig. 7, is shown for comparison (solid line).
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the electrolyte. In this case, the first law energy conversion efficiency (mCO) is equivalent to the ratio of the rate at which enthalpy is stored in CO to the rate of electrical energy used to produce it: mCO =
fCODH uCO (EW − EC)
(9)
where fCO is the rate of formation of the CO and DH uCO is the standard molar enthalpy change when CO is burnt in oxygen and EC is the potential of the counter electrode. The first law efficiency, calculated from Eq. (9) is shown in Fig. 10. An alternative expression for the overall energy conversion efficiency of the cell is the product of the Faradaic efficiency and the ideal enthalpic efficiency: mCO =
FCODH uCO nF(EW − EC)
(10)
where n is the number of electrons required in the electrochemical reduction of CO2 and F is Faraday’s constant. Eq. (10) is the inverse of the energy efficiency of a fuel cell [36]. As is shown in Fig. 10, the energy efficiency as calculated from Eq. (10) is in good agreement with that obtained using Eq. (9). At −1.5 V the energy storing efficiency was greater than 25% when calculated by both methods. According to the trend in Fig. 10, the maximum efficiency will be obtained as the applied potential approaches the onset potential. Eqs. (9) and (10) show that the energy storage efficiency can be increased by increasing either the Faradaic efficiency for CO formation or by decreasing the potential difference between the working and counter electrodes. Part of the energy lost during the formation of CO was observed as an increase in electrolyte temperature. Other losses include H2 and formate production at the cathode and hydroxyl and formate oxidation at the anode [37]. We measured the impedance spectrum of the eight cell unit, to estimate the energy lost due to electrode and electrolyte resistance. Before taking the measurements, the electrolyte was saturated with CO2 by passing the gas through the cells for several hours. The impedance spectra are shown in Fig. 11, along with the calculated frequency dispersion of an equivalent circuit. Using the notation of Boukamp [38], the equivalent circuit which gave a good fit to the measured values was a resistor (R1), a capacitor (C), an inductor (L) and a constant phase element (CPE: Q1, n1), all in series with a resistor (R2) and CPE (Q2, n2) in parallel. The values of the parameters of the equivalent circuit which gave a good fit to the impedance spectrum were R1 = 1.89 V, C=8.7× 10 − 4 F, Q1 = 2.1× 10 − 3, n1 = 0.81, L=1.6× 10 − 6 H, R2 = 0.35 V, Q2 = 6.6×10 − 3 and n2 = 0.86. From this analysis, the combined electrode and electrolyte resistance was 1.89 V. During CO2 reduction at − 1.5 V versus SCE, the potential difference between the working and counter electrode was
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Fig. 11. The magnitude ( z ) and the phase angle (h) of an impedance spectrum of the eight cell unit, obtained at an applied voltage of 5 mV rms. The electrolyte was 500 mM KHCO3 and the reference electrode was the counter electrode during this measurement. Also shown is z (solid) and h (dashed) from the calculated frequency dispersion of an equivalent circuit.
Fig. 12. Time dependence of current density during electroreduction of CO2 at a Au film cathode using a Pt anode in 500 mM KHCO3. At an applied potential of − 2.1 V, the gas flow rate was 8 ml s − 1 . At an applied potential of − 1.7 V, the gas flow rate was 33 ml s − 1
.
Under these conditions, the stability of this electrode was comparable to Dewulf and Bard’s SPE-supported GDE [19]. The resistance of the Au film remained unchanged during the experiment. The decline in current density may be due to one or a combination of factors, including poisoning of the electrode from impurities in the electrolyte or feed gas or from the reaction products. It is possible that CO may have been reduced to elemental carbon, lowering the number of available CO2 binding sites and resulting in a decrease in the current density. Alternatively, a trace metal in the KHCO3 such as Fe could have plated onto the Au electrode. If this occurred, the percentage of H2 produced would increase at the expense of CO [13]. Since adsorbed CO is expected to chemisorb at sites normally occupied by H2 [7], a decline in the current density could result from a decrease in H2 formation that would otherwise have occurred at sites occupied by CO. At an applied potential of − 1.7 V and a gas flow rate of 33 ml s − 1, the current density stabilized at around 8 mA cm − 2, remaining at this level for 48 h. After this time, the main reaction product was a gas other than CO. Based on the colour of the flame when this gas was burned in air, it was assumed to be H2. The formation of H2 supports the suggestion that metals like Fe may have plated onto the Au film. The decrease in current density over several days was accompanied by the formation of a blackish film, which was more pronounced near the exit of the electrode than at the entrance. This suggests the formation of a carbon deposit due to the reduction of CO. The formation of a carbon deposit on the electrode would probably also contribute to the observed decrease in the current density.
4. Conclusions 3.85 V and the total current flowing through the eight cell unit was −340 mA. Under these conditions, a potential drop of 646 mV between the working and counter electrodes was due to the resistance of the electrolyte and electrode, therefore a 5% improvement in the energy efficiency would be expected if the resistance of the electrode and electrolyte was eliminated.
3.7. Electrode stability In order to obtain an indication of the stability of the electrode over time, the current was measured as CO2 was passed over the series of eight Au electrodes. As shown in Fig. 12, at −2.1 V and a gas flow rate of 8 ml s − 1, the current rose to above 25 mA cm − 2. However, after 10 min the current density decreased by 40% over the following 100 min.
This work demonstrates the use of a porous thin Au film electrode for the electrochemical reduction of CO2 to CO, in aqueous KHCO3. The hydrophilic property of the supporting polymer substrate for the Au film was crucial for obtaining current densities comparable with those obtained at ordinary Au electrodes. Cyclic voltammetry in the presence and absence of CO2 confirmed that dissolved CO2 was reduced, rather than bicarbonate. The reaction kinetics were of fractional order in CO2 partial pressure, indicating that the rate-limiting step of CO2 reduction involves a bound species, possibly CO2−. However, the onset potential for CO2 reduction was about 900 mV more positive than the standard reduction potential for CO2− formation in aqueous media. Possible reasons for this were suggested.
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The Faradaic efficiency for CO2 reduction to CO was over 75%, which is comparable to that obtained at an Au-loaded GDE. The energy efficiency for energy storage in CO rose to a maximum of 35% near the onset potential for CO2 reduction. Impedance spectra of the electrodes indicated that up to a 5% increase in the energy efficiency could be expected if the electrode and electrolyte resistance was eliminated. Apart from heating and less than optimal Faradaic efficiency, the remaining energy losses are assumed to be due to the oxidation of COOH− and OH−. The current density during the electrochemical reduction of CO2 was stable for 48 h at an applied potential of − 1.7 V. At − 2.3 V, a steady decline in the current density was observed. Various reasons for the decrease in the current density were suggested.
Acknowledgements We thank Drs J.J. Lowke, D.L. Tilbrook and V. Braach-Maksvytis for helpful discussions and Drs R. Driver, C.P. Foley and A.F. Collings for critical reading of the manuscript. We also thank Drs I.C. Plumb, L.M. Besley and Mr M. Arnautovitch for assistance with gas analysis, Dr A. Molodyk for analysis of XRD spectra, Dr S. Lam for the SEM image and Mr G.R. Baxter for help in making the metal films.
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