Journal of CO2 Utilization 5 (2014) 53–59

Revisiting strontium-doped lanthanum cuprate perovskite for theelectrochemical reduction of CO2

Dimitri Mignard a,*, Rakesh C. Barik a,b, Arun S. Bharadwaj a, Colin L. Pritchard a, Marina Ragnoli c,Franco Cecconi c, Hamish Miller c, Lesley J. Yellowlees a

a The King’s Buildings, Mayfield Road, Edinburgh EH9 3JL, Scotland, UKb CSIR-Central Electrochemical Research Institute, Karaikudi, Tamil Nadu 630 006, Indiac Acta S.p.A., P.O. Box 143, 56021 Cascina (Pisa), Italy

A R T I C L E I N F O

Article history:

Received 25 June 2013

Received in revised form 23 December 2013

Accepted 27 December 2013

Keywords:

CO2 electrochemical reduction

Renewable power

Synthetic natural gas

Ethylene

Hydrogen

A B S T R A C T

The conversion of wind, solar or marine energies to fuels and chemicals can exploit the chemical

utilisation of carbon dioxide, and could lead to a more sustainable chemical industry. In particular,

electrolytic conversion of CO2 may suit the short-term variability of the energy supply while avoiding

issues associated with the large scale storage of hydrogen. Lanthanum cuprate perovskites were

described a while ago as promising electrocatalysts for the production of valuable higher alcohols,

ethylene and methane from CO2. However, confusion was apparent in the literature as to the nature of

the compound that was used, and the data on applied potentials was scarce. In this paper, we report on

the electroreduction of CO2 at a gas diffusion electrode loaded with a strontium-doped lanthanum

cuprate perovskite in 0.5 M KOH. The compound was prepared following the procedure that was

described in previous literature on CO2 electroreduction, and it was characterised as a mixture of bi-

layered La1.8Sr0.2CuO4 with tetragonal (La,Sr)CuO2.6. We also extended the range of temperatures (2–

40 8C) and pressures (1–43 barg). Methane and ethylene were major products of electroreduction, unlike

in previous work where alcohols were dominant. Results and the operational experience from this work

should help direct effort towards catalysts that have the advantages of copper metal without its

limitations.

� 2014 Elsevier Ltd. All rights reserved.

Contents lists available at ScienceDirect

Journal of CO2 Utilization

jo ur n al ho m ep ag e: www .e ls evier . c om / lo cat e/ jc o u

1. Production of chemicals from CO2 using renewable sourcesof power

‘Carbon capture’, which is the extraction of carbon dioxide (CO2)emitted by human activities, is widely seen as a key technology forreducing CO2 emissions from fixed sources [1–3]. As an alternative tostorage following capture, the chemical utilisation of CO2 taken fromindustrial sources [4–7] might allow the diversion of anthropogenicCO2 away from the atmosphere as a side benefit to more directlyproductive activities. This approach may also displace a moreexpensive, scarcier or ‘‘dirtier’’ feedstock by using CO2 instead.

One product that seems particularly attractive to prepare fromCO2 is ethylene, an essential precursor for the production of manypolymers and one of the organic compounds with the biggestworld production at 120 Mt/yr. It is produced by the cracking oflight hydrocarbons [8], a process that results in a substantialcarbon footprint at 1.1 kg CO2/kg ethylene, i.e. 0.7 mol CO2/mol ethylene [9]. There is currently substantial research effort

* Corresponding author. Tel.: +44 131 651 9024; fax: +44 131 650 6554.

E-mail address: d.mignard@ed.ac.uk (D. Mignard).

2212-9820/$ – see front matter � 2014 Elsevier Ltd. All rights reserved.

http://dx.doi.org/10.1016/j.jcou.2013.12.006

directed at the preparation of organic carbonates and polycarbo-nates by reacting CO2 with olefins, including ethylene [10].

A process for the production of ethylene from CO2 would also behelpful as an intermediate step towards ethanol synthesis, with theadvantage that the potential world demand for a fuel like ethanol ison a scale that approaches that of CO2 emissions. Typically,producing a fuel from CO2 requires the production of hydrogen byelectrolysis of water using a carbon neutral energy source likewind, solar or marine power [4], followed by a synthesis stepwhere the CO2 is reacted with hydrogen. However, using theserenewable energies produces a supply of hydrogen that varies withtime, which may jeopardise the synthesis step [5]. Thus, if onewished to manufacture either ethylene or ethanol from renewablepower and CO2, one must address the operation of the synthesisreactor at variable flow rates which involves issues of control,selectivity to products, and safety.

2. Electroreduction of CO2 on copper-based electrodes

These difficulties make it desirable to use electrolytic processesthat can accommodate the variability in time of the resources,while synthesising carbon-based fuels in one step. One such

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–5954

system that has been widely studied in the literature, due to itsunique ability to produce hydrocarbons or alcohols directly fromCO2 and water, is based on copper metal or its oxides as theelectrocatalyst, which was initially discovered in the first half ofthe 1980s by Hori et al. [11]. Following two decades of research inthe field, Gattrell et al. published a fairly recent survey on the topicof electrochemical reduction (‘electroreduction’) of CO2 at coppermetal electrodes [12].

Regarding copper as an electrocatalyst for CO2 reduction, thefollowing results stand out:

� It is well established that steady electroreduction of CO2 tomethane and ethylene with some hydrogen is achievable at coppermetal electrodes without noticeable degradation of activity [13],although prior purification of the electrolyte by pre-electrolysis isrequired and the length of experiments reported in the literature istwo hours at most. In potassium phosphate buffer bubbled withCO2 at the surface of polycrystalline copper [14], products appearin the following order as the potential is decreased to morenegative values: CO and HCOO� rise from �1.04 to 1.14 V vs.standard calomel electrode (SCE), reaching a maximum at around�1.44 to �1.49 V before decreasing; C2H4 appears at �1.34 V andCH4 at �1.44 V, and their rise corresponds with the decline of COand HCOO�. At �1.69 V vs. SCE and current densities of 5 mA/cm2,typical Faradaic efficiencies are ca. 44% for CH4 and 23% for C2H4,2% for CO, with the remainder about equally shared betweenHCOOH and H2 [14]. Similar results would be observed withhydrogen carbonate buffer [15].� The energy efficiency for the electroreduction of CO2 at copper

electrodes has been evaluated at 30–40% [16], as measured bythe ratio of chemical energy stored within the product to allenergy inputs to the process. While this is less than methanolsynthesis from CO2 plus hydrogen from water electrolysis (50–51% [4]) it seems a good starting point for catalyst and processoptimisation.� Unlike copper metal electrodes, copper oxide electrodes are able

to electrochemically reduce CO2 to methanol. The result was firstreported by Frese two decades ago [17], and again more recentlyby Chang et al. [18], but it had long been uncertain due to theinfluence of other factors (impurities from the electrolyte, andthe lack of long term stability of the oxide when exposed to areducing potential) until a thorough investigation by Le et al. [19]could confirm it. However, it seems that the copper oxides areprogressively reduced to copper metal and lose activity for theproduction of methanol [19].

In addition, most of the research that was reported in this fieldwas conducted at the surface of flat metallic copper electrodes,whether polycrystalline or single crystal plane electrodes. In thismode of operation, the electrolyte was usually saturated with CO2,with the effect of a limiting rate of supply of CO2 to the electrode forcurrent densities greater than a few mA/cm2. This was an issue,since high current densities of the order of several 100’s mA/cm2

would be required for practical applications. However, currentdensities were increased to this range of values either through theapplication of high pressure (for example [20], 400 mA/cm2 at apotential of �1.74 V vs. standard calomel electrode (SCE) and29 barg on a copper wire working electrode in 0.1 M KHCO3); orthe use of a gas diffusion electrode. In the latter case, the CO2 is fedthrough a porous working electrode made of conducting blackcarbon on which the copper metal or copper compound electro-catalyst is finely dispersed. With copper metal as the electro-catalyst and a Nafion1 membrane as the electrolyte, one study [21]reported current densities of 300 mA/cm2 at a potential of �2.69vs. SCE and atmospheric pressure. Combining pressurisation andgas diffusion electrode, another study [22] reported a current

density of 900 mA/cm2 at a potential of �1.31 V vs. SCE and 19 bargin 0.5 M KHCO3.

3. Aims of this work

In this paper, we report on the electrochemical reduction of CO2

at a gas diffusion electrode loaded with a strontium-dopedlanthanum cuprate perovskite. The experiments were conductedin 0.5 M KOH over a range of pressures and temperatures. Anelectrochemical cell was also built, to enable the use of the gasdiffusion electrode under high pressure.

This same system of electrocatalyst, electrode structure andelectrolyte was seemingly investigated by Schwartz et al. atambient pressure and temperature for the production of higheralcohols [23]. We noted that this result was unique in the literatureregarding the use of this electrocatalyst, and the high Faradaicyield of ethanol (30.7%) and n-propanol (10.0%) that were achievedunder the reported conditions (at 25 8C, 180 mA/cm2 in 0.5 M KOH,after 1.28 h). These authors had selected a lanthanum cuprate inthe hope that the surface planes of copper oxide in its crystaldomains might withstand reduction to copper metal (unlike purecopper oxides), thus allowing the stable production of alcoholsfrom CO2. They reported that selectivity to alcohols was markedlyenhanced, and that no leaching of copper ions from the gasdiffusion electrode in the electrolyte could be detected.

However, we noted that the nature of the lanthanum cupratecompound that was used in that earlier study was in fact unclear,and perhaps different from La1.8Sr0.2CuO4. This uncertainty arosebecause the very same results from that study were reportedelsewhere by the same team [24], but this time they were attributedto the closely related compound La0.9Sr0.1CuO3 instead. To the best ofour knowledge, no attempts at clarifying these results and theproperties of this system have been reported subsequently. It wasfelt worthwhile to clarify this confusion, because the resultsobtained with this type of copper oxide could inform ourunderstanding of the way that copper metal electroreduces CO2

and in what way the presence of the oxide and the other componentschanges the reaction pathways and the selectivity.

These same papers also reported that the production ofethylene and methane increased significantly when the tempera-ture was decreased to 2 8C. Therefore, there existed the possibilityof directly producing ethylene simply by changing the operatingconditions. Effects of pressure could not be investigated with theset-up used by these authors, which was a modified H-cell made ofglass, but would perhaps offer another means of improving theselectivity towards products in the condensed phase. Finally, therewas a scarcity of information regarding the potentials applied atthe working electrodes. It was mentioned that potentials at least asnegative as �2.2 V vs. standard calomel electrode (SCE) wererequired for production of alcohols, and that voltages were in therange �2.3 to �3.0 V.

In the light of these prior results, we attempted to clarify thenature of the lanthanum cuprate compound that had been madeback then, and to explore further the influence of potential,temperature and pressure on the selectivity to products. To achievethis, we made use of an electrochemical cell that had been designedfor the use of gas-diffusion-electrodes over a range of pressures andtemperatures. We chose 0.5 M KOH as the electrolyte in order toenable comparison with those earlier reports.

4. Materials and methods

4.1. Preparation of strontium-doped lanthanum cuprate perovskite

The strontium-doped lanthanum cuprate perovskite used inthis work was prepared by Acta S.p.A. following the description

Fig. 1. (a) Electrolytic cell and (b) magnified section of insert tube with supported

gas diffusion electrode.

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–59 55

given in [23], except that the nitrate salts of the metals used weredirectly bought from suppliers of chemicals rather than made byourselves from the corresponding oxides. This detail was deemedof no consequence for the composition of the final compound,which was expected to be the same as the one used in [23,24]. Asolution of citric acid (CAS 77-92-9; 5.76 g, Sigma–Aldrich 99.5%)in deionised water (40 mL) was added drop wise to a solution ofLa(NO3)3�xH2O (CAS 100587-94-8; 5.84 g on a dry basis, purity99.9%, Aldrich), Sr(NO3)2�H2O (CAS 152343-39-0; 0.46 g, Riedel-de-Haen, ‘extrapure’) and Cu(NO3)2�2.5H2O (CAS 19004-19-4;2.32 g, Aldrich 99.999%) in deionised water (60 mL). The solutionwas stirred for 30 min at room temperature. Most of the water wasthen removed under reduced pressure at 90 8C. The resulting oilymaterial was further dried under vacuum at 120 8C overnight. Theresulting solid was milled to a fine powder, transferred to a quartzreactor and heated to 980 8C for 12 h. We would expect this solid tobe the same as that obtained by the previously cited authors,except in the impurity content due to the suppliers of the reagentsbeing different. The solid thus obtained was milled again beforebeing weighed and dispersed in acetone (CAS 67-64-1; 85 mL/gcatalyst). The resulting dispersion was ultrasonicated for 10 min.Carbon black (CAS 1333-86-4; Vulcan VXC72R, kindly donated byCabot Inc.) was then added to the dispersion so as to obtain a 1:1proportion by weight of the catalyst with respect to the carbon;and the dispersion was left under agitation for at least 8 h beforethe acetone was evaporated under a vacuum using a rotavapour.

4.2. Preparation of gas diffusion electrodes

The catalyst powder was ground to fine particles using a mortarand pestle. The catalyst powder was weighed in a precision balanceand 40 wt.% PTFE liquid suspension in water (PTFE 2893, kindlydonated by Norkem Ltd.) was added into it to get a paste such that,after removing the water, the Teflon loading would be 40 wt.%. Afew mL of water was added dropwise while mixing with a stirringrod until the Teflon coagulated. The resulting paste was thenspread over a 21 mm diameter disc of teflonised carbon paper(supplier – Fuel Cell Scientific, 40% Teflon, thickness 190 mm) usinga spatula and a glass rod. This disc played the role of gas diffusionlayer as well as a fairly rigid backing. The electrode was dried (atroom temperature) and the paste roll-pressed further by using theglass rod to give a good attachment of the catalyst paste on to thesurface of the carbon paper and to prevent cracks. The electrodewas weighed and secured in place onto the working electrodesupport as explained in the next section.

4.3. Electrolytic cell

An electrolytic cell was built from a 1 in. OD stainless steelSwagelok1 cross, as shown in Fig. 1a. This choice was made inorder to be able to work at pressures up to that of the critical pointof CO2 at 73 barg. At one branch of the cross an insert thatsupported the working electrode was fitted, the design of which isdetailed in Fig. 1b. A stainless steel flanged end insert whosefunction was to transmit the current to the flat working electrodewas resting on another flanged insert made of PTFE for electricalinsulation from the rest of the rig, and had an electric lead solderedto it that came out at the other end of the insert.

A gas diffusion electrode (GDE) was placed onto a stainless steelmesh of the same diameter, with its teflonised carbon paper sidefacing the mesh, and the GDE and mesh were firmly held againstthe flanged stainless steel tip of the insert by the wrapping roundthe edge of some Parafilm1, with the electrolytically active layerfacing out as shown in Fig. 1b. The elasticity and adherence of theParafilm1 meant that the electrode was tightly held in place, andthe relative chemical inertness was deemed to allow little if any

interference with the processes at the electrode. To be morecertain, PTFE tape was wrapped on top to cover the Parafilm1, so asto minimise any influence of the Parafilm1, and the resultingexposed surface area of the working electrode was approximately1 cm2.

The counter electrode was a platinum foil made to face theworking electrode at a distance of 30 mm. The reference electrodewas a platinum wire, which was connected via an enamelledcopper wire to a calomel reference electrode outside the cell. Therestricted internal volume of the cell, combined with the need toprovide pressure tight ports for all connections to the electrodes,prevented the insertion of a conventional reference electrode or aLuggin capillary connected to a reference electrode. Therefore, weused platinum and copper wires to bring the potential at thesurface of the working electrode to the outside of the cell, where itcould be conveniently measured at the surface of a platinum wirethat was dipping inside a saturated KCl solution at roomtemperature (20 8C). This approach assumed that the contactpotentials between the copper wire and the platinum wires wouldbe negligible, and that the potentials between the platinum wiresand the electrolytes (inside and outside of the cell) would also benegligible (and in any case fairly constant in spite of the differenttemperatures and changing compositions within the electrolytes).While we have not found this method described elsewhere, it wasdeemed to be preferable to simply using a pseudo-referenceelectrode, which would not allow for the precise value of thepotential to be known.

The insert supporting the GDE was connected at the other endto another 1 in. Swagelok1 cross, allowing entry ports for the CO2

To ve nt

Thermo-

stated

bath

Liqu id

elec troly te

Cathode

lead

Ano de

lead Pressure

sensor lea d

Thermocou ple,

gas phase

Thermocou ple,

to liqui d in cell

Ref. elec trode

lead

CO2

Drier

Pressure

equi libratio n brid ge

Gas-

Diffus ion

Electrod e

Press ure Relief Valv e

with manual over ride

Nee dle va lve

To GC sampli ng

Fig. 2. Overall set-up for the cell. The shaded area represents the electrolyte,

contained in the bottom right Swagelok cross.

Hayesep/

PoropakMolecula r

sie ve

Hayesep/

PoropakMolecula r

sie ve

Before breakthro ugh of CO2 throu gh Hayesep/Porapa k column:

After brea kthrou gh of CO2 through Hayesep/Po ropak colu mn:

Fig. 3. Separation of gases in gas chromatograph; direction of flow and switching of

columns.

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–5956

feed and cathode collector lead. The two 1 in. crosses werethemselves connected at the top to a pair of ½ in. Swagelok1

crosses connected to each other by a tube whose function was toequilibrate the pressure between the two sides of the electrode.This bridge was provided with a view to avoid the dislodging orrupturing of the electrode should there be a pressure build up onone side of the electrode. The parts of the rig which were exposedto electrolyte were coated in PTFE on the inside. The outlet line,two thermocouples and a pressure sensor were also connected tothe ½ in. crosses. The assembly was kept in a thermostatic bath tocontrol the temperature between 2 and 40 8C. The overall set-up isillustrated in Fig. 2.

4.4. Analysis of products

Gas products were analysed using a Perkin Elmer 8700 GasChromatograph fitted with

� 10 port gas sampling valve, pneumatically actuated by relayswitches.� Argon carrier gas.� Two flame ionisation detectors (FID). One of the FIDs is fitted

with a methaniser, which is there to convert CO and CO2 tomethane for detection by the FID. The other FID is used for themanual injection port when performing liquid sample analysis.� A TCD detector for detecting hydrogen and oxygen. Unfortu-

nately, this detector suffered progressive degradation and finallydestruction on two occasions, which was traced to theprogressive obstruction of a flow restrictor (the cause of whichwas unclear). Therefore, we were unable to quantify hydrogenand oxygen gas products with certainty.� A stainless steel Hayesep/Porapak column (3 m � 1/800 OD, DIP

80–100 mesh), in series with molecular sieve column (80–100mesh of 13�, adsorbing molecules smaller than 9 A). The firstcolumn interacted strongly with CO2 and as a result it preventedthe large signal associated with CO2 from masking the signalsfrom other compounds. The light gases were then separated inthe molecular sieve. Finally, the order of the columns wasswitched just before the CO2 had finished crossing the Porapakcolumn (Fig. 3).� Liquid products including alcohols and aldehydes were analysed

by transferring 5 mL of the electrolyte solution to a 20 mL vial,which was immediately sealed with a septum cap and held at60 8C for 20 min, thus allowing equilibrium compositionsbetween liquid and gas phase to be established. 500 mL samplestaken from the headspace of the vials with a syringe through theseptum cap were then manually injected into a BP20 (WAX)capillary column from SGE (25 m � 0.53 mm i.d., film 1 mm) that

was fitted in the same GC as the columns used for gas analysisand products detected by FID. Organic acids including formicacid were separated in an 8.5 � 300 ORH HPLC column suppliedby Speck&Burke Analytical, and detected by UV spectroscopywith a Perkin-Elmer UV/VIS LC290 spectrophotometer.

4.5. Electrolysis

A 0.5 M potassium hydroxyde solution was prepared using99.99% KOH (Sigma–Aldrich) and deionised water (15 MV). Usinga Parstat 2275 potentiostat (Princeton Research Applied), theelectrolyte was pre-electrolysed at 0.1 mA for 24 h using asacrificial working electrode which was prepared in the sameway as the working electrode to be used in the actual experiment.It was hoped that this procedure would scavenge from theelectrolyte the contaminants to which the working electrode wasspecifically sensitive. This preliminary step was conducted in abeaker. Meanwhile, the whole rig was assembled with theworking electrode fitted to it, except that the top of theelectrolysis cell was left open. The rig was filled with carbondioxide at atmospheric pressure, then the pre-electrolysedelectrolyte was transferred to the cell and the top fitting put inplace to seal the cell. Electrolysis at constant current in the range200–300 mA/cm2 was then conducted at specified temperatureand pressure (in the range 2–40 8C and from 1 barg to saturatedvapour pressure of the CO2 at the given temperature), withsamples of gas and liquid taken at times t = 20 min, 60 min and120 min after the start of the electrolysis. When sampling, thewhole of the liquid in the cell was removed and replaced withfresh, pre-electrolysed electrolyte.

5. Results

Results reported here were obtained at a current density ofeither 200 or 300 mA/cm2, with an extensive study of productdistribution over the range of temperatures and pressuresavailable with the rig. No liquid products were detected with thiscatalyst. Although the hydrogen product could not be quantifiedreliably, it was detected when the TCD was still working. Theresults presented here are for gases other than hydrogen.

The set of conditions that were actually measured was 2 8C,10 8C, 20 8C and 40 8C at 1 barg and at 10 barg; 2 8C at 33 barg; and10 8C at 43 barg.

Fig. 4 plots the current efficiency to methane production after20 min vs. the compensated potential applied to the workingelectrode with respect to the standard calomel electrode (itselfwith a potential of +0.244 V vs. SHE at 25 8C; sets of differentpressures and temperatures are distinguished). Fig. 5 does likewisefor ethylene and Fig. 6 for carbon monoxide.

Fig. 4. Current efficiencies to methane and potentials with respect to standard

calomel electrode at different pressures and temperatures for La1.8Sr0.2CuO4

catalyst, after 20 min.

Fig. 5. Current efficiencies to ethylene and potentials with respect to standard

calomel electrode at different pressures and temperatures for La1.8Sr0.2CuO4

catalyst, after 20 min.

Fig. 6. Current efficiencies to carbon monoxide and potentials with respect to

standard calomel electrode for La1.8Sr0.2CuO4 catalyst after 20 min.

Fig. 7. Change with time of the current efficiency to methane and potential with

respect to standard calomel electrode, at 10 8C, 10 barg and 300 mA/cm2.

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–59 57

By current efficiency to a given product is meant the ratio of thenumber of electrons that contribute towards the synthesis of thatproduct from CO2, divided by the total number of electrons fed tothe electrode during the duration of the electrolysis so far.

Most of the experiments showed that activity towards CO2

reduction decreased during the course of the electrolysis. Fig. 7illustrates this behaviour, describing the evolution with time of thecurrent efficiency with respect to methane production at 300 mA/cm2, 10 barg and 10 8C. The current efficiency dropped from 10.1%in the first 20 min to 3.4% between 20 and 60 min, and then to 0.3%between 60 min and 120 min. Fig. 7 also includes a plot againsttime of the potential at the working electrode with respect to the

calomel reference electrode, showing that it decreased from�2.23 V to �2.74 V between 20 min and 60 min, but stayed at thatlevel or recovered very slightly afterwards to reach �2.67 V at120 min.

6. Discussion

6.1. Products from CO2 electroreduction

We detected only methane, ethylene and carbon monoxide asproducts from CO2 electroreduction. While we cannot exclude thatother products might have been made in comparable amounts, thisis not likely since we could not detect oxygenates in the liquid (ifthey had formed, we would have expected to detect at least tracesof these even if they were being destroyed at the anode), and therewas no sign of ethane amongst the products (making it unlikelythat there were C3+ hydrocarbons). Therefore, we concluded thatCO, methane and ethylene were the only major products from CO2

reduction. Some of these gases may have reoxidized at the anode,but this may not have occurred to a significant extent given thatthis anode was submerged in the electrolyte. In conclusion, webelieve that the sum of their current efficiencies should be close tothe total current efficiency of CO2 reduction. Overall, this meansthat the balance of current efficiencies to 100% could have beenapproximated by the hydrogen product, which was detected,though not quantifiably.

Figs. 4–6 all refer to experiments conducted at either 200 or300 mA/cm2. Initially, this choice was made on the basis ofensuring a practical operating current density from the point ofview of industrial application (a high current density is required inorder to keep the size of industrial electrolysers down, and hencecurb the associated capital cost). However, Figs. 4–6 show that thehighest current efficiencies were mostly found at the least negativeapplied potentials. This suggests that both the geometry orstructure of the electrode, and the electrolyte concentration orits nature, were far from optimal, and that less negative voltagesmight have resulted in still higher current efficiencies.

Conditions for which maximum current efficiencies wereobtained depended on the product under consideration and thepotential at the electrode. As reported in Fig. 4, the currentefficiency to methane after 20 min reached a maximum of 20% at10 barg, 2 8C and potential of �1.35 V vs. SCE, current density of300 mA/cm2. (Under the same conditions, current efficiencies toethylene and carbon monoxide were 5.3% and 0.2%, respectively.)By contrast, the maximum current efficiency to ethylene after20 min of electrolysis was 9.4%, achieved at 1 barg, 20 8C and�1.94 V vs. SCE, 200 mA/cm2 (in these conditions, current

0

200

400

600

800

1000

1200

1400

1600

1800

20 25 30 35 40 45 50 55 60 65 70

Degree s, 2 theta

No

rmali

zed

in

ten

sit

y

*

*

*

*

* *

****

*

**** *o

o

o oo oo

o

o *

Fig. 8. XRD spectrum (Cu K-a) of the strontium-doped lanthanum cuprate used in

this work. ‘*’ symbols denote tetragonal (La,Sr)CuO4 (including, La1.8Sr0.2CuO4; ICSD

65847, but possibly with a different ratio of Sr with respect to La) and ‘o’ symbols

denote tetragonal (La,Sr)CuO2.6 (for example, La0.84Sr0.16CuO2.6; ICSD 73885).

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–5958

efficiencies to methane and carbon monoxide were 11.4% and 2.2%,respectively), while the maximum current efficiency to CO was6.8%, achieved at 1 barg, 2 8C and �1.5 vs. SCE, 300 mA/cm2 (withmethane at 13.8% and ethylene at 5.4%).

6.2. Comparison with earlier work

Generally, the results do not look comparable with earlierreports [23,24], since with the same electrolyte (0.5 M KOH) theirauthors observed the production of alcohols rather than hydro-carbons, and in our experiments the activity could not bemaintained over two hours. This was in spite of applying potentials(�1.35 to �3.80 V vs. SCE) and current densities (200–300 mA/cm2) that were comparable to theirs (�2.3 to �3.0 V and 180 mA/cm2, respectively).

Looking at the detail of our procedure, it can be seen that theelectrolyte was not circulated, and that throughout the experimentit remained in contact with the CO2 in the headspace provided atone end of the pressure equilibration bridge. While this set-uphelped us explore the effect of pressure, it could have created rapidsaturation with CO2 and subsequent reaction with hydroxyl ions.The resulting drop in pH will have favoured the formation ofhydrocarbons rather than alcohols. This is in contrast with earlierpublished work [23,24] where a larger volume of electrolyte wasrecirculated, and the only contact with CO2 was through the poresof the electrode where reduction would have directly competedwith neutralisation.

Another effect of electrolyte recirculation in that earlier workmight have been to create sufficient agitation in the vicinity of theelectrode surface to affect the ratio of adsorbed CO and hydrogenon the electrocatalyst. While this phenomenon was demonstratedto occur at the surface of flat copper electrodes [20], it was CO thatwas preferentially desorbed (a fact that was readily explained bythe small value of its enthalpy of adsorption on copper), and somore reduced compounds like hydrocarbons and alcohols werefavoured in the absence of agitation when CO was allowed toremain adsorbed at the surface for long enough for furtherreduction to proceed. However, the data was not sufficient toindicate whether alcohols rather than hydrocarbons would befavoured at some intermediate level of agitation.

Another possibility is the effect of impurities, with our pre-electrolysis procedure possibly failing to remove some of themetallic ions that were initially present in the electrolyte. Inaddition, an accurate reproduction of the earlier work would haverequired the use of exactly the same reagents and electrolytes aswell as the same finer details of the experimental procedures, apoint that was demonstrated very convincingly [13] by a series ofstrictly identical experiments of CO2 electrolysis at copperelectrodes but each time using KHCO3 electrolytes from differentsuppliers.

In fact, the exact composition of the catalyst might have been astrong determinant of the outcome. It is now clear that the catalystthat was used in this work, and which was prepared following thesame procedure as described in the earlier work [23,24], was amixture of two phases rather than just either La1.8Sr0.2CuO4 orLa0.9Sr0.1CuO3. Fig. 8 presents an XRD spectrum of the strontium-doped lanthanum cuprate used in this work. In the region 2u = 30–368, the positions of the peaks match those published by Schwartzet al. [23], but the height ratios are quite different, especially the(0 2 0) reflection at 2u = 338, which appears much higher for theActa catalyst. In particular, the height ratio of the 338 peak to the33.68 peak is about 1.5 in Fig. 8, and 0.33 in [23]. However, Fig. 8and, with hindsight, also the narrow portion of XRD spectrum thatwas reported in [23], both fit a superposition of two spectra: thatfor tetragonal (La,Sr)CuO4 (including, La1.8Sr0.2CuO4; ICSD 65847,but possibly with a different ratio of La and Sr) and that for

tetragonal (La,Sr)CuO2.6 (for example, La0.84Sr0.16CuO2.6; ICSD73885), the latter giving the 338 peak. A similar observation ofthese two phases being present in strontium-doped lanthanumcuprate that was prepared in similar conditions was reported in[25]. It could be that the ratio of these two phases is critical for theselectivity towards either alcohol or hydrocarbons during CO2

electroreduction.

6.3. Loss of activity

The apparent loss of activity when comparing results at 20 min,60 min and 120 min as exemplified in Fig. 7 could be explained byany of the following:

� Systematic failure by the pre-electrolysis process to thoroughlyremove the impurities; perhaps slow leakage of contaminantsfrom the cell wall (for example, the removal of the sacrificialelectrode required switching off the potentiostat, as required bysafe operating practice. However, the corresponding change inpotential might have had the effect of re-dissolving theimpurities).� Leakage of the electrolyte through the gas diffusion electrode and

its contact with the walls of the stainless steel tube thatsupported the electrode (this was often observed to haveoccurred during the experiment).� Slow conversion of CO2 to carbonate ions in the KOH solution will

lead to substantial losses in conductivity [26]. This would haveincreased the voltage applied to the cell (since the currentdensity was maintained constant), with the result that theworking electrode would have been exposed to an increasinglynegative potential that favoured hydrogen evolution rather thanCO2 reduction, and hence an apparent loss of activity. Althoughthe electrolyte was always renewed when samples were beingtaken, the pores inside the catalytically active layer of theelectrode could not be washed. Therefore, the observation inFig. 7 of the applied potential becoming more negative between20 min and 60 min is consistent with this expectation. The slightrise to more positive potentials between 60 min and 120 minwas not observed in all experiments (it typically kept decreas-ing), and it might be explained by accidental changes to themacrostructure of the electrode, for example following thedetachment of fragments leading to a smoother surface andhence easier detachment of bubbles, which would decrease theoverall resistance to the passage of ions between the electrodeand the electrolyte (discussed in the next bullet point).

D. Mignard et al. / Journal of CO2 Utilization 5 (2014) 53–59 59

� On several occasions, disintegration of the gas diffusion electrodecoating was seen to have occurred during the experiment, withblack debris of the electrode coating found within the electrolytewhen reopening the cell after the experiment. At the same time,the surface coating of the electrode would look visibly degraded.These fragments were detached from the electrode by thevigorous gas evolution.

6.4. Further requirements for practical application

From the point of view of commercial application, the yields ofethylene remained relatively low at 9% current efficiency. Energyefficiency is also kept low by the cathodic potentials in ourexperiments being more negative than �1.3 V vs. SCE. This wouldbe made worse by the relatively low conductivity of the electrolyte(0.5 M KOH) when compared with commercial electrolytes thatare used in water electrolysis (concentrated KOH or Nafion1). Thispoor energy efficiency seems to be a characteristic of CO2

electroreduction using copper and its oxides in aqueous electro-lytes, seemingly supported by recent DFT modelling results [27],and suggesting that other directions be explored; and yet theunique ability of copper to electrochemically convert CO2 toethylene with good yields is compelling as a starting point for abetter catalyst. Given the strong contribution of solvation energy tothe activation barrier of CO2 for the first electron transfer, and thedetrimental effect of low electrolyte conductivity to the overallenergy efficiency of the cell, a better choice of electrolyte would bea first step. Ogura published work in which the electrolytes usedwere fairly concentrated solutions of potassium halides in therange 2–3 M [28]. In addition, Cook et al. found that CO2 reductionto methane, ethylene and ethane could be conducted at potentialsless negative than �0.67 V vs. SCE, using polymer membraneelectrolytes in a cell with the following configuration: CO2/Cu/Nafion1/Pt, N2 (90%), H2 (10%) [29]. These results suggest thatrelatively efficient CO2 electroreduction might be achieved bycopper-based electrocatalysts if appropriate cell configuration andelectrolyte environment can be found.

Finally, hydrogen was likely to be a major product (although wecould not quantify it). However, in view of the conclusions reachedby a large study conducted by a consortium of gas and energycompanies and European research institutions from 2004 to 2009,the possibility exists of injecting this hydrogen together with themethane product into a natural gas pipeline. The NaturalHy projectconcluded that in most cases the existing transmission anddistribution networks for natural gas should be able to handle up to20% hydrogen safely. Unmodified existing domestic gas applianceswere operable on the mixture, provided that they complied withthe latest standards; industrial boilers, gas engines and otherapplications might need evaluating on a case by case basis [30].

7. Conclusions and further work

The compound prepared in earlier work for which the resultswere reported in two publications [23,24] was neither pureLa1.8Sr0.2CuO4, nor pure La0.9Sr0.1CuO3, but in fact a mixture oftetragonal (La,Sr)CuO4 and tetragonal (La,Sr)CuO2.6, thus clarifyingthe confusion between the data from these two sources. However,we did not observe the production of alcohols at the La1.8Sr0.2CuO4

gas diffusion electrode in 0.5 M KOH unlike what was reported in[23,24]. Instead we observed the production of methane, ethylene

and carbon monoxide. There appeared to be an optimumtemperature and pressure of 10 barg, 2 8C for the production ofmethane with current efficiency of 20%; and 1 barg, 20 8C for theproduction of ethylene with current efficiency of 9%. Even thoughwe followed the same preparation procedure, our electrocatalystappears to have contained a larger proportion of the (La,Sr)CuO2.6

phase, which might explain why alcohols were not observed asreduction product. Other experimental factors that might havecontributed to enhance the selectivity towards hydrocarbonsrather than oxygenates could have been the free surface of theelectrolyte in contact with CO2, or the use of reagents andelectrolyte from different suppliers, or the lack of recirculation ofthe electrolyte. Future work should focus on appropriate choice ofelectrolyte that may allow increased energy efficiency throughhigher conductivity, and on lowering the overpotentials at thesurface of the electrocatalyst.

Acknowledgements

The authors would like to express their gratitude to Acta S.p.A.and the Carbon Trust for funding this work, and to Dr Ronald Brownfor hosting the experiment in his laboratory at the School ofChemistry at the University of Edinburgh.

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