indian Journal of Chemistry Vol. 41A, March 2002, pp. 472-477

Influence of pH and supporting electrolyte on electrochemical reduction of CO2 using nickel(II) macrocyclic complex of 1, 3, 6, 9, 11, 14 -

hexaazacyclohexadecane as catalyst at HMDE

M Aulice Scibioh t, B Viswanathan+* & V R Vijayaraghavan t tOepartme nt o f Physical Chemistry , University o f Madras, Guindy Campus,Chennai 600025 , Ind ia

+Oepartment of Chemistry, Indi an Institute of Technology Madras Chennai 600036, India

Received 9 January 2002; revised 30 Novdelllber 2001

Electrochemical studies o n the titl e compound using cyclic voltammogram (CV) and control potential electro lys is (CPE) techniques reveal that it reduces CO2 electrocatalytically at -1.36V /SCE at hang ing mercury drop e lec trode (HMDE) in aqueous mcdium using LiCI04 as a supporting electrolyte. The products are found to be CO and Hz in 3:2 mo le ratio in thc gaseous phase as detec ted using gas chromatography (GC) and trace amounts of formic acid in soluti on phase as detec ted using colorimetric technique . The I / Id values (where h is the kinetic c urrent measured in the presence o f CO2 and Id is the diffusion c urrent measured in N2 atmosphere) o!Jserved at various pH values show that pH 5.0 is best suited for CO2 reductio n. In add ition, the hydrophobicity/hydrophilic ity near the electrode surface provided by the cati on of the supporting salt and it s influence on CO2 reducti on is di scussed .

In recent years the reduction of carbon dioxide (C02) has attracted considerable interest, largely aimed at finding processes for the conversion of CO2 either into useful organic chemicals or into useful fuel (CO)I. The main issue in these studies is the activation of the quite inert CO2 molecule. Direct electrochemical reduction , which is the most exploited approach, requires a large negative overvoltage on the majority of electrode materials studied2-8. In order to reduce this overvoltage, some alternative approaches have been proposed, which includes the electrocatalytic reduction of CO2 using transition metal complexes as the mediator of the electron transfer. With this aim,

I f h . 9-1 1 . 12 h ' 13 comp exes 0 r enlUm , Iron , rut enlUm , cobale 4 , rhodium l3.15 and nickeI1 4. 16.17 with N or P containing ligands were used either in a homogeneous phase'i· lo.12.13.15- 18 or by suitably attaching them to the electrode surface (modified electrodes) 1 1,1 4.19. The results obtained using a nickel(ll) azamacrocyclic complex of I , 3, 9, I I, 14-hexaazacyclohexadecane for the reduction of CO2 at -1.36Y /SCE at hanging mercury drop electrode (HMDE) in aqueous medium using LiCI04 as a supporting electrolyte are reported elsewhere2o. The products were found to be CO and H2 in 3:2 mole ratio in the gaseous phase as detected using gas chromatography (GC) and trace amounts of fo rmic ac id in solution phase using colorimetric

technique. Since the pH of the medium has been observed to influence the feasibility of the reaction as well as the nature of the products formed, study of the effect of pH on the electrocatalytic reduction of CO2 is one of the objectives of the present work . In addition, the role of supporting electrolyte is also studied, since the supporting electrolyte is not regarded only as a component to give ionic conductivity but the cations of the support ing electrolytes play a significant role in product di stribution by changes caused in the double layer structures . Therefore, the resultant hydrophobicity/-hydrophilicity near the electrode provided by the cation of the supporting electro lyte on the electrocatalytic reduction of CO2 is also considered. For choosing and performing electrochemical experiments in the presence of various electro lytes, reference was made to a summary of some important properties of widely used solvents and supporting electrolytes21 along with their potenti al windows,

. 22 1] proVIded by Sawyer and Roberts and Mann -. .

Materials and Methods Aqueous solutions were made from analytical

grade salts and doubly di stilled water. Potassium hydrogen phthalate buffer was used to study the redox reactions at pH 4.0 and 5.0, potass ium dihydrogen

SCIBIOH et al. : ELECTROCHEMICAL REDUCTION OF CO2 473

phosphate buffer was used for pH 6.5 and tris(hydro-xymethyl)aminomethane buffer and disodium hydrogen phosphate buffer were used for pH 9.0 and 12.0 respectively. Other near pHs such as 10.5, 5.5 , and 5.0 were obtained by the gradual addition of HC104• Purified N2 and CO2 were used for deaeration.

The solvents used are acetonitrile (CH3CN), N, N'-dimethyl formamide (DMF) and N, N'- dimethyl sulphoxide (DMSO) and of spectroscopic grade from Qualigens, India. The supporting salts such as LiCI04 (Aldrich), NaCI04 (Aldrich), KCI04 (Fluka), NH4CI04 (Aldrich), TBAP (Fluka) and TEAP (Fluka) were of spectroscopic grade and used as purchased.

Gas samples (500 Ill) were drawn at various time intervals with a gas tight syringe through a septum and were analyzed in a gas chromatograph provided with a thermal conductivity detector (TCD) and 20% carbowax on chromosorb column using nitrogen as carrier gas . The detecting limits for CO and H2 are 0.25% and 0.012% respectively. Formic acid was determined by a colorimetric method24.

Electroc/,elllicall1leasuremellts Cyclic voltammetric (CY) experiments were

performed with a Wenking ST 72 potentiostat and voltage scan generator with a Graphtec XY recorder WX 2300, using a three electrode system consisting of a locally fabricated hanging mercury drop (HMDE: 0.0310 cm2) working electrode, mercury pool auxiliary electrode and saturated calomel (SCE) reference electrode, brought into the main compartment via luggin capillary. Aqueous (0.1 M) LiCI04 was used as supporting electrolyte. All solutions were purged with N2 (or CO2) where required, for 0.5 h prior to each scan. In general, cycli~ voltammograms were recorded from -0.9 to -1.6 Y at scan rates in the range 25-400 mYs- J • Experiments were run at 298±2 K.

A two-compartment gas-tight electrolysis cell was used for the controlled-potential electrolysis (CPE) experiments . The working electrode was mercury pool (I 1.65 cm2). The counter electrode Pt was separated from the main compartment by GO sinter. The electrolytic solution (15 ml) was deaerated by bubbling CO2 (or N2) for 30 min and el ectrolyzed at -1.60 Y /SCE for 5h.

Results and Discussion Role of pH

. The reduction of CO2 requires protons under neutral conditions. Hence, the measurement of pH in

the presence and absence of CO2 in aqueous solution of the title compound [NiL]2+ (L = I , 3, 9, 11, 14-hexaazacyclohexadecane) would indicate a possible reaction route for the reduction of CO2. It is observed that the pH of the aqueous solution of complex dropped from 7.0 in the absence to 5.0 in the presence of CO2. This has prompted us to study the electrocatalytic reduction of CO2 under acidic as well as alka line conditions.

The electrochemical studies have been carried out using complex [NiL]2+ of 0 .5 /JIM concentration in 0.1 M LiCI04 at various pH viz., 4.0, 5.0, 6.5 , 9.0 and 12.0 and cyclic voltammograms are recorded under N2(a) and CO2(b) atmospheres . When cyclic voltammograms are recorded at pH 4 .00, the pH was changed from 4 .00 to 3.89 in the presence of N2 and CO2. No peaks were observed either under N2 or CO2 atmosphere. This might be due to the fact that in acidic pH, the reduction of H20 (or W) competes effectively with the reduction of CO2. Therefore, thi s observation may be accounted for on the basis of a highW/C02 ratio or decomposition of the electrocatalyst. In the later case, free nickel would be produced catalyzing the formation of H}5.

Typical cyclic voltammograms of [NiLf+ at pH 5.0 under N2(a) and CO2(b) atmospheres are shown in Fig. I. The Mp was found to be 60 mY in N2 atmosphere. The pH of the solution was 5.0 and 4.87 ± 0.02 under N2 and CO2 respectively. Studies under pH 6.5 using potassium dihydrogen phosphate buffer show results similar to that in unbuffered solutions2o, i.e. in 100% H)O.

The values of Illd (where h is the kinetic current measured from the amplitude of the signal in the presence of CO2 and Id is the diffusion current measured in the same way in N2 atmosphere) for the complex at pH 5.0, 6.5, 9 .0 and 12.0 were found to be 5.2,3.6,3.4 and 2.8 respectively . It was found that h measured at pH 5.0 was higher than that measured at any other pH. This might be due to the presence of optimum concentration of H+ for CO2 reduction. At lower pH, the H+ reduction would be favored and at higher pHs, [H+] supply would become insufficient for CO2 reduction. Hence pH 5.0 should be an ideal one for the effici ent catalytic reduction of CO2. The Manichaen character of CO2 imposes two problems: On the one hand CO2 competes with hydrogen for electrons, and on the other hand, most of the reduction reactions require protons. The Illd values were found to decrease as the pH increases from 9.0 to 12.0. This

474 INDIAN J CHEM., SEC. A, MARCH 2002

is due to the lower availability of H+ ions for the reduction of CO2. The Jp,c values measured in CO2 atmosphere were plotted against pH for [NiLf+ and is shown in Fig, 2. Since, Ip,c value is maximum at pH 5.0, it was further confirmed that the pH 5.0 is the most suitable one for the reduction of CO2 in aqueous medium.

Role of supporting electrolytes The electrolyte medium, which consists of solvent

and the supporting electrolyte, exerts a major influence on the nature of the electrochemical process, The chemical properties of the electrolyte medium affect the electrochemical reaction mechanism in the same way solvents affect the normal reaction chemistry26, From the earlier literature27 , it is understood that the supporting electrolytes influence both solubility and mechanism of reaction in the case of reduction of CO2, Also, the effect of supporting salt has not been studied extensively; the supporting electrolytes have been regarded as components to give the necessary ionic conductivity. For example, Murata and Hori 28

-F ---

~ T c 5 !J A '-'- 1 :J

U

-0·9

\

\ '<

\

\

\

\ \ \ \ \

\ \ , ,

-1 .3

\ " \ /-.\ \ ,_/ \ \ b \ ,

\ \ \ \

\ \ \ I \ \ I, 1\ 1\

" II

)

-1·7

E,V vs SeE

Fig I-Cyclic voltammogram of [NiLf+ (0.5 mmol dm·3) at HMDE in phthalate buffer (PH 5.00); 0, I mmol dm·3 LiCl04 under N2 (a) and CO2 (b), Scan rate = 100 mVs·

l•

reported the effect of alkali metal cations on the product distribution caused by changes in the double-layer structure28, and Bockris et al.29•3 1 reported the effect of tetraalkylammonium ions, Saeki et al. 32

studied the role of supporting electrolyte in CO2 + methanol system and discussed the alteration of hydrophilicity/hydrophobicity near the e lectrode surface provided by the cation of the supporting electrolyte and its effect on CO2 reduction. Hence, it is proposed to study the electrochemistry of the title complex in the presence of CO2 using vari ous electrolytes, Table 1 represents the va lues of I / Id for various supporting electrolytes studied 111 the desirable solvent compositions,

Role of tetrabutylammollium perchlorate (TEAP) Since, 0,1 M TBAP is not soluble in 100% H20,

electrochemical studies could not be performed in 100% aqueous solution of the complex [NiLJ"+ , The insolubility of TBAP in water i attributed to the bulkiness of butyl group, Even though, TBAP is completely soluble in 100% CH3CN, 100% DMSO and in 100% DMF, the reduction of CO2 is not observed in dry non-aqueous solvents (in the presence of complex), as the reduction of CO2 requires protons. Therefore, aqueous mixtures of the solvents were taken for studies, TBAP(O, I M) in {he absence of the complex [NiL]2+ (for all solvent mix tures studi ed) showed no peaks in the potential range -0,9 to -1.7 Y, both in N2 and in CO2 atmospheres, The solvent mixtures used for the electrochemical reduction of CO2 in the presence of complex using TBAP are: 80% CH3CN 20% H20 (v/v), 80% DMF 20% H20 (v/v)

20

O~ __________ L _ _________ L-________ ~ o 5 10 15

pH

Fig 2-Plot of Ip.c vs pH for [NiLf+ (0.5 mmol dm·3) under and CO2 atmosphere. Scan rate = 100 mVs·l .

SCIBIOH et al.: ELECTROCHEMICAL REDUCTION OF CO2 475

Table I-Studies on the effect of supporting electrolyte on current for the reduction of CO2

Supporting Solvent Electrolyte , O.IM

200 TBAP 80%CH)CN 5.65

80% DMF 6.75 70%DMSO 4.3

TEAP 100% CH)CN 3.4 80%CH)CN 6.45 70% DMF 7.0

KCI04 100% H2O 2.6 70%DMSO 3.0

NaCI04 100% H2O 2.55 80% DMF 2.8

and 70% DMSO 30% H20 (v/v). (In each case, further increase in water content causes turbidity due to the poor solubility of TBAP). The IIId values for the above mentioned solvent mixtures in the presence of 1.0 M TBAP and 0.1 M [NiLf+ are given in Table 1. It is noted that the hlld ratio is higher for 80% DMF-20% H20 (v/v) solvent mixtures in the presence of 0.1 M TBAP. The hlld values for all solvent mixtures studied in 0 .1 M TBAP were found to be higher than the hlld values observed using 0.1 M LiCI04• It is inferred that TBAP suppresses H2 evolution more effectively than does LiCI04 .

Role of tetraethylammonium perchlorate (TEA?) Tetraethylammonium perchlorate is readily and

completely soluble in water at room temperature. Therefore, e lectrochemical studies were carried out using 0.1 mM complex in 1.0 M TEAP both in N2 and CO2 atmosphere. TEAP is completely soluble also in DMF and in CH3CN, but no catalytic reduction of CO2 takes place in the presence of [NiL]2+ in dry non-aqueous solvents, as the reduction of CO2 requires protons. Hence, aqueous mixtures of DMF and acetonitrile were used. The IIId values for the experiments in aqueous and aqueous solvent mixtures are given in Table 1. All solvent compositions (non-aqueous : water = v/v) were hied and among the soluble ranges, those solvent compositions which gave higher IIId values alone are given in Table 1. The hlld ratio for 70% DMF-30% H20 (v/v) mjxture is the highest among the solvent mixtures studied.

Role of potassium perchlorate (KCL04) Since potassium perchlorate is completely soluble

in water at room temperature, studies have been conducted using 0.1 M [NiL]2+ in 1.0 M KCl04 both in

Scan rate, v(mVs·1 )

100 5.7 6.8 4.4 3.5 6.5 7.2 2.7 3.1 2.6 2.9

50 7.5 8.0 6.5 4.2 8.3 8.5 2.9 4.0 2.85 3.8

Table 2- The hi!" values at various concentrations of [NiLf+ ill 100% H20

Complex, hi!" mmol.dm·) 0.25 6.20 0.50 3.60 0.75 3.03 1.00 2.95 1.25 2.86 1.50 2.53 1.75 2.57 2.00 2.65

N2 and CO2 atmospheres. KCI04 is insoluble in acetonitrile and slightly soluble in 80% and 60% CH3CN (v/v). Therefore, studies could not be performed in CH3CN and in their aqueous mixtures. Though KCI04 is fully soluble in 100% DMSO there is no catalytic reduction of CO2 found in this medium and hence aqueous mixtures were chosen. Amon,g the various compositions tried, 70% DMSO-30% H20 (v/v) is found to exhibit the highest IIId value. The hlld values obtained for the solvents studied using 0.1 M KCI04 are given in Table 1. It is seen that the value of IIId ratio is higher in 70% DMSO-30% H20 (v/v) mixture than in 100% H20. This may be due to the presence of excess of H+ ions (in H20), which are competing with CO2 for electrons. By comparing the results of studies in TBAP, TEAP and KCI04 , hydrogen evolution is greater (as inferred from li ld values) when KCI04 is used as electrolyte.

Role of sodium perchlorate (NaCL04) NaCI04 is fully soluble in H20 and in' DMF.

Therefore, electrochemical studies have been conducted in 100% H20 but since there is no catalytic reduction of CO2 in dry DMF, studies were carried

476 INDIAN J CHEM ., SEC. A, MARCH 2002

out for aqueous mixtures of DMF and 80% DMF-20% H20 (v/v) mixture is found to give the highest Ii/Id value. The Ii/Id values for the solvents studied in the presence of NaCl04 are presented in Table 1. Here, the value of Ii/Id ratio is higher in 80% DMF-20% H20 (v/v) than in 100% H20 due to the presence of excess of H+ ions which compete with CO2 for electrons. The trend observed for NaCl04 is similar to that observed for KCl04 .

Role of ammonium perchlorate (NH4CI04) Since ammonium perchlorate is fully soluble in

100% H20 , studies have been performed in aqueous medium in the presence of 0.1 M complex. NH4Cl04 is not fully soluble in DMSO, but aqueous DMSO mixtures were chosen for the study. In both media, hydrogen evolution is highly predominant.

Role of lithium perchlorate (LiCL04) The results of [NiLf+ using 1.0 M LiCl04 have

already been presented and discussed2o. However, it is worth mentioning a few important observations for the purpose of comparison. LiCl04 is readily and completely soluble in water at room temperature. Table 2 represents the hlId values for various concentrations of the complex in aqueous medium. The Ii/Id value is 2.95 for 0.1 M complex of [NiL]2+ in 100% H20 at 100 mVs-

l. Table 3 represents the

variation of hlId values for the complex in various solvent compositions using lithium perchlorate as supporting electrolyte. Eventhough LiCl04 is soluble

in CH3CN, DMF and DMSO, there is no catalytic current found in dry non-aqueous media under CO2 atmosphere. However, aqueous mixtures of the solvents electrocatalyse CO2 in the presence of complex. Among various solvent mixtures studied, the Ii/Id values for the more ideal mixture at v = 100 mVs- 1 are as follows: 80% H20-20% DMF (v/v) = 3.7, 80% H20-20% DMSO (v/v) := 3.7 and 20% CH3CN-80% H20 (v/v) = 2.8.

Considering all the results , the following conclusions can be drawn concerning the effect of pH and supporting electrolytes on the catalytic reduction of CO2 using [NiL]2+ as a catalyst.

(i) In aqueous medium, pH 5.0 is the ideal one for the effective reduction of CO2.

(ii) The hlId values are higher in TEAP than in TBAP in 80% CH3CN. Also, the hlId values are higher in TEAP in 70% DMF than in TBAP using 80% DMF (only the most efficient composition was chosen in each case). Therefore, it is concluded that tetraethylammonium perchlorate suppresses hydrogen evolution more than tetrabutylammonium perchlorate. This observation is similar to that of the studies of Yoneyama et al. 32, where the effect of electrolytes on current efficiency for CO2 reduction at semiconductor electrodes and product distribution were investigated in aqueous media.

(iii) The Ii/Ill values for the electrolytes are in the following order (0.1 M electrolyte at scan rate 100 mVs- 1) in 100% H20 medium: TEAP > LiCl04 > KCl04 > NaCI04 . (As TBAP is insoluble in H20 , it is

Table 3-Studies of scan rate variation in solvents of various compositions. [NiLf+ = 1.0 mmol.dm·3 ; [LiCI041 = 1.0 mmol.dm·3; Temp.= 298±2 K

Solvent

DMF

DMSO

50 60 70 80 90 100 20 40 60 80 100 20 40 60 80 100

400

3.2 3.35 3.5

3.45 3.40 3.0

3.25 3.32 3.40 3.45 3.00 1.9 2.2 2.3 2.9 3.0

200

3.2 3.4 3.6 3.5

3.45 2.92 3.30 3.35 3.45 3.50 2.92 2.43 2.50 2.55 2.85 2.92

hlld Scan rate, V (mVs-l)

100 50

3.25 3.45 3.5 3.7 3.8 4.1 3.7 4.2 3.5 3.9

2.95 3.1 3.25 3.60 3.40 3.70 3.65 3.90 3.70 4.0 2.95 3.1 2.5 4.2 2.6 5.45

2.65 6.7 2.8 7.0

2.95 3.1

SCIBIOH et al.: ELECTROCHEMICAL REDUCTION OF CO2 477

not included for comparison). The organic perchlorates suppress hydrogen

evolution more effectively than does alkali metal perchlorates.

(iv) For aqueous mixtures of organic solvents the l/ld values follow the order: TEAP > TBAP > LiCI04 > KCI04 > NaCI04'

The hydrophobic atmosphere at the electrode provided by the TBA ion (or TEA ion) may be favourable for CO2 reduction. Whereas, hydrogen evolution proceeds efficiently in the hydrophilic atmosphere provided by lithium or other alkali cations. We noticed that hydrophilicity and hydrophobicity provided by the supporting electrolyte in this study, whereas they were provided by the electrode surface in the system studied by Y oneyama etal 33 .

(v) Ammonium ion in NH4CI04 also provides a hydrophilic atmosphere like lithium or other alkali ion. But the hydrogen evolution was predominant while using NH4CI04 and hence, it favors hydrogen reduction over CO2 reduction.

References I The organic and bio-organic chemistry of carbon dioxide,

edited by Inone I S & Yamazaki M (Halstead Press, New York) 1982.

2 Haynes L V & Sawyer D T, Anal Chem, 39 (1967) 332. 3 Amature C & Saveant J M, JAm chem Soc, 103 (1981) 5021. 4 Eggins B R & McNeill J, J eLectroanal Chem, 148 (1981)

5021. 5 Bradley M G, Tysak T, Graves D J & Vlachopoulas N A, J

chem Soc Chem Commun, (1983) 349. 6 Koudelka M, Mannier A & Augustynsky J, J electrochem

Soc, 131 (1984) 745. 7 Vassiliev Y B, Bagotzky V S, Osterova N V, Khazova 0 A &

Mayorova N A, J Electroanal Chem, 189 (1985) 271. 8 Vassiliev Y B, Bagotzky V S, Khazova 0 A, & Mayorova N

A, J electroanal Chem, 189 (1985) 295. 9 Hawecker J, Lehn J M & Ziessel R, J chem Soc Chem

Commun, (1984) 328.

10 Sullivan B P, Bolinger C M, Conrad D, Vining W J, Murray R W & Meyer T J, J chem Soc Chem Commun, (1985) 1414.

11 O'Toole T R, Margerum L D, Westmoreland T D, Vining W J, Murray R W & Meyer T J, J chem Soc Chem CommulI, (1985) 1416.

12 Tezuka M, Yajima T, Tsuchiya A, Matsumoto Y, Uchida Y & Hidai M, JAm chem Soc, 104 (1982), 6834.

13 Bolinger C M, Sullivan B P, Conrad K, Gilbert J A, Story N & Meyer T J, J chem Soc Chem Commun, (1985) 796.

14 Meshitsuka S, Ichikawa M & Tamaru K, J chem Soc Chem Comlllun, (1974) 158.

15 Slater S & Wagenknecht J H, J Am chem Soc, 106 (1984) 5367.

16 Fisher B & Eisenberg R, J Alii chelll Soc, 102 ( 1980) 7361 . 17 Beley M, Collin J P, Ruppert R & Sauvage J P, J chem Soc

Chem Commlln, (1984) 1315. 18 Keene F R, Creutz C & Sutin N, Coord Chem Rev, 64 (1985 )

247. 19 Lieber C M & Lewis N S, JAm chem Soc, 106, (1984) 5033. 20 Aulice Scibioh M & Vijayaraghavan V R, Bull Electrochem,

13(6), (1997) 275. 21 Murata A & Hori Y, Bull ~hem Soc Japan , 64 (1991) 123. 22 Schaap W & McKinney D,'Experimental electrochemistry for

chemists, (Wiley, New York). 1974, Chapter 4. 23 Mann C K, Non-aqueous soLvents for electrochemical lise

edited by A J Bard, Vol. 3, (Dekker, New York), 1968. pp. 57-134.

24 Grant W M, Anal Chem, 36(A), (1964) 267. 25 Beley M, Collin J P, Ruppert R & Sauvage J P, J Am chem

Soc, 108 (1986) 7561. 26 Brown E R & Sandifer J R in Physical methods of chemistry,

edited by R W Roositer and J F Hamilton , (Wiley, New York), 2nd Ed., 1986, Vol. 2, Ch. 4, pp. 273.

27 Murphy 0 J & Srinivasan S, Electrochemistry in transition, (Plenum Press, New York), 1992. Chapter 22. Blajeni B A. Electrochemical reduction of carbon dioxide.

28 Murata A & Hori Y, Bull Chem Soc Jpn, 64 (1991) 123. 29 Taniguchi I, A-Blajeni B & Bockris J 0' M, J electroanal

Chem, 161 (1983) 179. 30 Bockris J 0' M & Wass J C, J Electrochem Soc, 136 (1989)

2521. 31 Taniguchi I, A-Blajeni B & Bockris J O' M, J electroanal

.chem, 161 (1983) 385. 32 Saeki T, Hashimoto K, Kimura N, Ornata K & Fujishima A. J

electroanal Chem, 390 (1995) 77. • 33 Yoneyama H, Sugimara K & Kuwabata S, J electroanal

Chem Interfacial Electrochem, 249 (1988) 143.