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Electrocatalytic reduction of CO2 by nickel(II) cyclam
199
J. Electroanal. Chem., 292 (1990) 199-215 Elsevier Sequoia S.A., Lausanne
Electrocatalytic
reduction of CO, by nickel(I1) cyclam
Study of the reduction mechanism on mercury by cyclic voltammetry, polarography and electrocapillarity Masamichi Fujihira *, Yoshiki Hirata and Kosaku Suga Department of Biomolecular Yokohama 227 (Japan) (Received
29 March
Engineering,
Tokyo Institute
of Technology,
4259 Nagatsuta,
Midori-ku,
1990; in revised form 9 May 1990)
ABSTRACT The electrocatalytic reduction of CO1 on mercury with Ni(II)-cyclam was studied in detail by cyclic voltammetry, polarography and electrocapillarity of aqueous solutions with and without the catalyst under N,, CO and COa. It is concluded from these electrochemical measurements that (i) the adsorbed complex (Ni(I)-cyclam,,) on Hg, not the reduced complex in solution Ni(I)-cyclam, is the active catalyst for the CO* reduction; (ii) both Ni(II)- and Ni(I)-cyclam can adsorb on Hg, and the oxidation states of the adsorbed complexes depend on the electrode potentials; (iii) Ni(I)-cyclam adsorbs on Hg even at - 2.0 V (vs. SCE) under N,, but the surface concentration of Ni(I)-cyclam,, decreases gradually with a decrease in the potential and desorbs completely at potentials more negative than - 1.7 V under CO or under CO or CO, is due to the formation of COa; and (iv) the desorption of Ni(I)-cyclam,, unadsorbable Ni(I)-cyclam.CO by the reaction of Ni(I)-cyclam in solution with dissolved CO or electrocatalytically generated CO from COa.
INTRODUCTION
The electrochemical reduction of CO, on metallic cathodes requires typically negative potentials more than- -2.0 V vs. SHE (- 2.24 V vs. SCE) [l]. Energy consumption is reduced by the use of molecular electrocatalysts [2] such as metallophthalocyanines [3-S], metalloporphyrins [9,10] and various tetraazamacrocyclic complexes [ll]. CO or formate ion was shown to be the major product of the reduction [5,6,11]. Several macrocyclic complexes have also been used for the photoelectochemical reduction of CO, to CO with various semiconductors at rela-
l
To whom correspondence
0022-0728/90/$03.50
should
be addressed.
0 1990 - Elsevier Sequoia
S.A.
200
tively low negative potentials [12,13]. Various acyclic complexes have also been proposed as electrocatalysts, including an iron-sulphur cluster [14], and complexes of rhenium [15], rhodium [16,17], ruthenium [17] and nickel [18]. Sauvage and co-workers have reported the exceptional efficiency and selectivity of nickel(I1) cyclam (cyclam is an abbreviation for 1,4,8,11-tetraazacyclotetradecane) as an electrocatalyst for the reduction of CO, on Hg in water [19,20]. The catalytic reduction involves the production of CO by the following electrochemical reaction [1,2]: CO,+2H++2e-+CO+H,O
E”‘=
-0.76V(vs.SCE)
(I)
This catalyst has also been used for the photoelectrochemical [21,22] and photochemical [23] reduction of CO,. However, the details of the electrochemical reduction mechanism of CO, by Ni(II)-cyclam have not yet been clarified. In particular, present knowledge of catalytically active species that are adsorbed on Hg is limited. Therefore, in this work the role of the adsorbed catalyst on the electrochemical reduction of CO, on Hg was studied in detail by cyclic voltammetry (CV) [24], polarography [25] and electrocapillarity [25] in an aqueous Ni(II)-cyclam solution under various atmospheric conditions (N,, CO and CO,). As a reference, the cyclic voltammetric behaviour on glassy carbon (GC) was also studied. EXPERIMENTAL
Cyclam was received from Aldrich. N-Tetramethylated cyclam (1,4,8,11_tetramethyl-1,4,8,11-tetraazacyclotetradecane, abbreviated as TMC) was prepared as described by Ciampolini et al. [26,27]. The nickel(I1) complexes with these ligands were prepared and obtained as perchlorate according to the literature [28]. The following gases (Tomoe) were used without further purification: nitrogen (99.9995%), carbon monoxide (99.9%) and carbon dioxide (99.99%). Aqueous solutions were prepared from analytical grade reagents and pure water from a Milli-Q system (Millipore Ltd.). The pH of the solution under CO, was adjusted by adding NaHCO, to aqueous 0.1 M NaClO,. Under N, or CO, an aqueous solution of 0.1 M NaClO, without buffer was used as the electrolyte. CV, polarography and electrocapillarity were carried out with a Nikko Keisoku potentiostat by using an electrochemical cell with a salt bridge containing aqueous 0.1 M NaClO, in order to remove chloride contamination from the saturated calomel electrode (SCE) used as the reference electrode. An amalgamated gold (Hg-Au) disk (1.5 mm in diameter), a CC disk (3.0 mm in diameter) or a dropping mercury electrode (DME) was used as the working electrode together with a GC rod counter-electrode. The Hg-Au disk electrode was prepared by polishing a Au disk electrode with emery paper (No. 2000) and then 0.3 pm alumina, sonicating it with Milli-Q water and finally amalgamating it by dipping the electrode into GR-grade Hg. The GC disk electrode was also polished and sonicated similarly prior to use. All measurements were done at room temperature (ca. 20°C). All the potentials given below are referred to the SCE.
201 RESULTS
AND
DISCUSSION
Cyclic voltammetry
of Ni(II)-cyclam
itself in water under N, on GC and Hg-Au
The cyclic voltammograms of Ni(II)-cyclam in an aqueous solution containing 0.1 M NaClO, under N, on the GC and Hg-Au disk electrodes are shown in Figs. la and lb, respectively. Two reversible waves at + 0.80 and - 1.56 V, corresponding to the Ni(III)/Ni(II)and Ni(II)/Ni(I)-cyclam redox couples respectively, were observed on GC; while only one reversible wave at - 1.56 V for the Ni(II)/Ni(I)cyclam couple could be recorded on Hg-Au because of Hg dissolution at positive potentials. The sweep rate (v) dependence of the peak heights of these reversible waves, i.e. a linear dependence of the peak currents on z?/*, indicates clearly their diffusion-controlled characteristics [24]. The height of the reoxidation peak of the Ni(II)/Ni(I)-cyclam redox couple was slightly decreased at slow sweep rates,
10pA
(b)
(cl
0.8jJA I
-2.0
-1.0 E
0 /
v
vs.
1.0 SCE
Fig. 1. Cyclic voltammograms of 1 mM Ni(II)-cyclam aqueous solution containing 0.1 M NaCIO, under N, on (a) a GC disk electrode and (b) a Hg-Au disk electrode. (c) Sweep rate dependence of the adsorption prewave of the same solution on the same Hg-Au disk electrode. Sweep rate/V s-l: (a, b) 0.1; (c) 0.1, 0.2, 0.3, 0.4, 0.5.
202
probably [29]:
because of the catalytic
2 Ni(I)-cyclam++
nature of Ni(I)-cyclam
2 H,O + 2 Ni(II)-cyclam2++
and/or the limited lifetime of Ni(I)-cyclam Ni(I)-cyclam++
for hydrogen evolution
H, + 2 OH-
(2)
due to a ligand loss reaction [30]:
2 H,O + Ni(I),‘, + cyclamH:+
+ 2 OH-
(3)
The reduced overpotential for hydrogen evolution upon repeated potential cycling was more apparent on GC than on Hg-Au. The deposition of nickel on GC by the electrochemical reduction of Ni(I)& generated in eqn. (3) may account for the activation of hydrogen evolution. It is also interesting to note that one small reversible prewave of the Ni(II)/Ni(I)-cyclam redox couple described previously by Sauvage and co-workers [20] was observed clearly only on Hg-Au. Figure lc shows the sweep rate dependence of the prewave observed at ca. - 1.33 V on Hg-Au in 1 mM Ni(II)-cyclam aqueous solution. The almost linear dependence of the peak height on the sweep rate indicates that this wave can be ascribed to the adsorbed Ni(II)/Ni(I)-cyclam (Ni(II)/Ni(I)-cyclam.,) redox couple on a Hg-Au electrode surface [24]. The positive shift of the adsorption peak potential compared with the half-wave potential of the redox couple in solution at -1.56 V shows clearly that Ni(I)-cyclam adsorbs more strongly on Hg than does Ni(II)-cyclam. The importance of Ni(I)cyclam., in the catalytic reduction of CO, will be described later. Cyclic voltammetry
of Ni(II)-cyclam
itself in water under CO on GC and Hg-Au
In order to clarify the electrocatalytic reduction mechanism of CO,, the cyclic voltammetric behaviour of Ni(II)-cyclam under CO was examined prior to CV under CO,. CO was found to be the main product of CO, electrolyses on Hg in the presence of Ni(II)-cyclam in water [20]. In Fig. 2a a cyclic voltammogram of 1 mM Ni(II)-cyclam in 0.1 M NaClO, under CO on GC between - 1.50 and + 1.15 V is shown. The potential sweep was started from - 1.0 V in the positive direction. During the first cycle (solid line) between -1.0 and +1.15 V, no change was observed between the cyclic voltammograms under N, and under CO. Namely, only one reversible wave was observed at the same potential of + 0.80 V as that observed under N, which was responsible for the Ni(III)/Ni(II)-cyclam redox couple without any further coordination reaction by CO. In contrast, the other reversible wave corresponding to the Ni(II)/Ni(I)-cyclam redox couple under N, was shifted under CO in the positive direction by 170 mV and its reoxidation peak height was further diminished. In the second positive sweep (dashed line), two new small peaks appeared at ca. - 0.6 and + 1.0 V. When the lower end of the potential sweep was extended to - 1.8 V as shown in Fig. 2b, another new anodic wave appeared at +0.6 V and its peak height increased with repeated potential cycling. The origin of the new peaks observed at - 0.6, + 0.6 and +l.O V has not yet been clarified, but Sauvage and co-workers [20] found that the
203
Fig. 2. Cyclic voltammograms of 1 mM Ni(II)-cyclam aqueous solution containing 0.1 M NaCIO, under CO on a CC disk electrode with (a) a narrow potential sweep and (b) a wide potential sweep, and (c) on a Hg-Au disk electrode with a narrow potential sweep. Sweep rate: 0.1 V s-l. The solid lines and dashed tines indicate the first and second cycle, respectively.
Ni(1) carbonyl complex described below is much less stable in H,O than in DMF. By taking account of a similar positive shift by CO observed for the more stable redox couple of Ni(II)/Ni(I)-TMC [30-321, the following coordination reaction of CO to electroproduced Ni(I)-cyclam is most likely responsible for the positive potential shift: Ni(II)-cyclam++
e-+
CO + Ni(I)-cyclam.
CO
(4)
Several tetraazamacrocyclic complexes of Ni(1) having carbonyl as the fifth ligand with square-pyramidal geometry have been generated in solution [33]. As no new peaks appeared for the CV of Ni(II)-TMC in the presence of CO, except for the potential shift of the reversible wave at -0.99 V under N, to -0.75 V under CO [32], the carbonyl formation in eqn. (4) cannot be responsible for the new anodic peaks. The same potential shift as for the Ni(II)/Ni(I)-cyclam redox couple under CO was also observed on Hg-Au as shown in Fig. 2c. When the sweep rate was changed from 100 to 500 mV s-l, there was almost no change in the shifted half-wave potential at -1.39 V or in the peak separation between the cathodic and anodic
204
peaks. Rather, an increase in the reoxidation wave-height of the redox wave accompanied by a relative decrease in the height of the anodic wave of a by-product at ca. -0.7 V was observed with the increase in sweep rate. These observations indicate that the coordination reaction in eqn. (4) is very rapid and fairly reversible. As is obvious from a comparison of Figs. 2a and 2c, the peak potential of the by-product was shifted from - 0.6 to - 0.7 V upon changing the electrode material from GC to Hg-Au. In addition to this anodic peak, another new sharp anodic peak at - 0.06 V and a new redox wave at + 0.05 V appeared. The peak currents of the latter waves were characteristic on Hg-Au and increased with repeated potential cycling or by extending the lower potential limit. As these waves were not observed on GC, a reaction between some other unknown by-products of Ni(II)-cyclam and Hg is most likely the cause of these new waves near 0 V on Hg-Au. CV of 1 mM free ligand cyclam on Hg-Au under N, gave a redox wave at - 0.07 V. This wave can be related to Hg(II)-cyclam/Hg’ redox couples [31,34]. Therefore, the redox waves observed at +0.05 V under CO in Fig. 2c suggest that, in the presence of CO, the ligand loss reaction in eqn. (3) is accelerated at negative potentials and the resulting free cyclam gives a redox wave for Hg(II)-cyclam/Hg’ which is shifted slightly to positive potentials by carbonyl formation.
Cyclic voltammetty
of Ni(II)-cyclam
in water under CO, on GC and Hg-Au
The cyclic voltammograms of 1 mM Ni(II)-cyclam under CO, in aqueous 0.1 M NaClO, containing 35 mM NaHCO, (pH 6.2) on GC and Hg-Au are shown in Figs. 3c and d, respectively. In Figs. 3a and 3b, their background currents observed in the same electrolyte under CO, without Ni(II)-cyclam (pH 6.2) are shown. On GC, two cathodic peaks are seen at - 1.5 and - 1.8 V in the negative sweep, while one cathodic peak at - 1.8 V and one small anddic peak at - 0.7 V are seen in the positive sweep. On Hg-Au, one large cathodic peak at -1.5 V in the negative sweep and two small anodic peaks at ca. -0.7 and -0.1 V in the positive sweep are seen. In comparison with their background currents and the cyclic voltammograms of Ni(II)-cyclam under N, and CO, the catalytic reduction of CO, in the presence of Ni(II)-cyclam is obvious. But the efficiency on GC was found to be much lower than that on Hg-Au in terms of the current density in A cme2, i.e. if we take account of the four times larger area of GC than that of Hg-Au. In addition to this magnitude of the catalytic current density, the onset potential for the catalytic current was much more favourable on Hg-Au than on GC. This effect of the substrate electrode material on the catalytic nature supports strongly the view that the adsorbed species plays an important role in the catalysis. The appearance of the anodic peaks under CO, in Figs. 3c and 3d are ascribable to the by-products, which were not observed under N, (Fig. 1) but were observed under CO (Fig. 2) and suggests that the main reduction product of CO, is CO on both electrode materials.
205
I
I
,
I
I
-2.0 E
/
I
-1.0 V
I
a
I
8
I
t
0 vs.
SCE
Fig. 3. Cyclic voltammograms of aqueous solutions containing 0.1 M NaClO, and 0.035 M NaHCO, (pH 6.2) under CO2 on a GC disk electrode (a) without and (c) with 1 mM Ni(II)-cyclam, and on a Hg-Au disk electrode (b) without and (d) with 1 mM Ni(II)-cyclam. Sweep rate: 0.1 V s-l.
Dependence
of the catalytic
currents on the concentration
of Ni(II)-cyclam
in solution
As is apparent from the comparison between the adsorption prewave of Ni(II)cyclam in Fig. lc and the catalytic reduction wave of CO, in Fig. 3d, the two onset potentials for the adsorption and the catalytic waves are in good agreement. This coincidence implies that Ni(I)-cyclam, on Hg is responsible for the catalysis, as Sauvage and co-workers first pointed out [20]. We also repeated CV on Hg-Au under CO, in aqueous solutions containing various concentrations of Ni(II)-cyclam. As shown in Fig. 4, we could reproduce an adsorption isotherm-like dependence of the peak intensity on the logarithmic concentration of Ni(II)-cyclam in M. Again, this observation supports strongly the importance of Ni(I)-cyclam., for the catalytic reduction of CO, in comparison with Ni(I)-cyclam formed in the aqueous bulk.
206
log [ NiW-cyclam
I
Fig. 4. Plot of the cathodic peak current density of CV for the catalytic electrode vs. the logarithmic concentration of Ni(II)-cyclam in M.
CO, reduction
on a Hg-Au
disk
From the cyclic voltammetric behaviour of Ni(II)-cyclam on Hg under N,, CO and CO, described so far, it is not clear whether Ni(II)-cyclam was adsorbed on Hg as well as Ni(I)-cyclam. The potential dependence of the amount of Ni(I)-cyclam., is also not obvious from CV. A study of the electrocapillarity was attempted, therefore, for Hg/aqueous 0.1 M NaClO, without Ni(II)-cyclam and Hgfaqueous 1 mM Ni(II)-cycIam in the supporting electrolyte systems under N,, CO and CO,. The electrocapillary curves were obtained approximately by measuring the average drop time for ten successive drops, keeping the applied electrode potential and the height of Hg column constant, and by repeating the drop time measurements at different potentials. The typical drop time curves thus obtained are shown in Fig. 5. In the absence of Ni(II)-cyclam, an approximately parabolic curve was obtained under CO which coincided with the curves under N, and COz, indicating no appreciable adsorption of CO on Hg. In contrast, in the presence of Ni(II)-cyclam a decrease in the drop time from that observed without the catalyst was seen under N, at potentials more negative than -0.3 V. The difference in drop time increased gradually with decreasing electrode potential until the potential reached the redox potential for Ni(II)/Ni(I)cyclam at - 1.56 V. Beyond this potential, the difference in drop time decreased with a decrease in the potential, but it was still lower than that observed without the catalyst even at - 2.0 V, It is interesting to note that a break in the curve was not
201
vv V
(a)
0
(b) (cl (dl
A 0 V
V V V %
V 0 %7 0
IQ
40
0
V V
0 “Ooo 0 3.0
%
Q
,,,,,,,,,,,,,,,,,,,P~ 0
-0.5 E
-1.5
-1.0 /
V
vs.
-2.0
SCE
Fig. 5. Drop time curves of aqueous 0.1 M NaCIO, solutions with 1 mM Ni(II)-cyclam under (a) N,, (b) CO and (c) CO2 with 0.035 M NaHCO, (pH 6.2), and (d) without Ni(II)-cyclam and 0.035 M NaHC4 under N,.
observed at the potential for Ni(II)/Ni(I)-cyclam.,, but at the potential for Ni(II)/Ni(I)-cyclam in solution. The decrease in drop time observed at all potentials more negative than - 0.3 V indicates clearly that both Ni(II)- and Ni(I)-cyclams can adsorb on Hg, although the redox states of the adsorbed complexes, of course, depend on the electrode potential. Under CO, a difference in the drop time with and without the catalyst was also observed at potentials less than -0.3 V, but a break in the curve was found at ca. -1.35 V which corresponded to the half-wave potential for the reduction of Ni(II)-cyclam in the presence of CO, i.e. the reaction described in eqn. (4). At potentials of ca. - 1.7 V or less, almost no difference in drop time with or without the catalyst was observed. The difference between the drop time curves under N, and CO implies that coordination of CO to the electrogenerated Ni(I)-cyclam in solution diminishes the amount of Ni(I)-cyclamad. Under CO,, the drop time curve was very similar to that observed under CO. The lack of adsorbability of Ni(I)-cyclam . CO produced by the reaction between Ni(I)cyclam in solution and the electrocatalytically generated CO may also account for the curve under CO,. Polarographic
study of the catalytic reduction of CO, on mercury
In Figs. 6a-6c are shown the polarograms in instantaneous currents of 1 mM Ni(II)-cyclam under N,, CO and CO,, respectively. Under N,, a reversible wave
208
(a)
fb)
III1
I
I
-2.0
I
Iii!
i
-1.5 E
Fig. 6. Polarograms
/
I
I
- 1.0 v
vs.
SCE
of 1 mM Ni(II)-cyclam
CO and (c) CO2 with 0.035 M NaHCO,
aqueous
solution
containing
0.1 M
NaClO.,under(a) NZ,(b)
(pH 6.2).
with a half-wave potential of -1.56 V can be seen, but an adsorption prewave, which is well known for methylene blue [25], is not obvious in Fig. 6a. The adsorption prewave for Ni(II)-cyclam could be barely seen under high magnification of the current axis. The reversible wave was shifted to a positive potential by carbonyl formation in the same way as in CV under CO (Fig. 6b). However, the polarographic wave of the catalytic reduction of CO, was complex, as in CV. Maximum currents were not necessarily observed at the end of growth of the DME. The outline connecting the maximum currents is similar to the cyclic voltammogram in Fig. 3d; that is, the decrease in the currents on the negative side of the peak potential of CO, in CV was also observed in Fig. 6c. This immediately suggests that the abruptly decreased current in CV in Fig. 3d was not due to the depletion of CO,, but due mainly to the decrease in catalytic activity of the Hg surface with Ni(I)-cyclam,. An excess of C&, however, is not enough for the rate to be regarded as a quasi-first-order reaction, SO that the
209 b +
a
d +
L !L jL f +
+
L
JL
5
JL j + ?f
;
: -L/J
I
~OJJA
h
‘2;. Fig. 7. Changes in the instantaneous current as a function polarography of 1 mM Ni(II)-cyclam aqueous solution containing under CO1 (pH 6.2). Applied electrode potential/V (vs. SCE): -1.40, (e) -1.45, (f) -1.50, (g) -1.55, (h) -1.60, (i) -1.65, - 1.85, (n) - 1.90, (0) - 1.95.
of time at various potentials in the 0.1 M NaClO, and 0.035 M NaHCO, (a) -1.25, (b) -1.30, (c) -1.35, (d) (j) -1.70, (k) -1.75, (1) -1.80, (m)
catalytic peak current at ca. - 1.4 V decreased slightly with the decrease in height of the Hg column indicating a small effect of depletion of CO*. It is worth noting that the potential ca. - 1.7 V where a sudden decrease in current was observed in the polarogram is in good agreement with the potential where adsorption of the catalyst ended under CO, as indicated by the drop time curve in Fig. 5. To study the polarogram in detail, the dependence of the instantaneous current on time, i.e. I-t curves, was observed as a function of the applied electrode potential. Typical I-t curves recorded from dropping of the preceding DME with a pen recorder are shown in Fig. 7. At the potentials between the onset of ca. - 1.15 V and - 1.40 V giving the maximum catalytic current, the instantaneous current (1) increased smoothly with time (t). Log I vs. log t plots in this potential region are shown in Fig. 8. At all potentials, the slopes were initially 2/3, which corresponds to the increase in the surface area of the DME with time [25]: A = 0 85m2/3t2/3
(5)
where m is the flow rate of mercury in g s-l. But the slopes can be seen to deviate from this theoretical slope afterwards, during the drop-life. This deviation started at an earlier stage of drop growth and became larger with an increase in the applied negative potential. The proportionality of the catalytic current to the surface area
210
Fig. 8. Log I-log t plots for changes in instantaneous currents as a function of time at various potentials in V vs. SCE: (a) -1.40, (b) - 1.35, (c) -1.30, (d) -1.25.
gives a good indication of the rapid adsorption equilibrium of the catalyst and the surface catalytic nature of the CO, reduction. Depletion of the CO, concentration in the vicinity of the electrode surface is most likely the cause of the deviation from the theoretical slope in this potential region. As is obvious from the I-t curves at further negative potentials in Fig. 7, the rise in the initial current became increasingly steeper, but the current levelled off and even fell in the last seconds of the drop-life. At - 1.70 V, a rapid rise and fall in I was observed in the initial 1 s. At further negative potentials, the currents were kept low during their drop-lives. At potentials more negative than -1.80 V, again a monotonic increase in the current with time was observed. The current observed in this most negative potential region is attributable to hydrogen evolution, described later. In Fig. 9, the instantaneous currents at 0.3, 0.5 and 3.0 s after the beginning of drop growth are plotted as a function of the applied electrode potential. From the potential dependence of the current observed at 0.3 s, the catalytic reduction of CO, was found to be accelerated by the increase in negative potential until the potential reached - 1.7 V, where the catalyst was desorbed in the presence of catalytically generated CO. If the potential were negative enough for the observed current to be
211
-200 -2R
E
/
-1.5 V vs.
-1.0 SCE
Fig. 9. Instantaneouscurrentsat (a) 0.3 s, (b) 0.5 s and (~1 3.0 s after the beginningof drop growth as a function of the electrode potentialplotted from the 1-f curves in Fig. 7.
diffusion controlled of CO,, the I-t curve would behave as predicted by the IlkoviE equation [25]: 1, = 0.732nF~~‘/~~~/~~‘/~
(6)
and the current for each growing drop would increase with time during the drop-life. Therefore the decrease in the instantaneous current with time observed at potentials between - 1.40 and - 1.70 V should be attributed not to CO, depletion, but to the gradual decrease in the surface catalyst concentration during the drop-life. The decrease with time in the amount of Ni(I)-cyclam=~ may be due to the increase in the concentration of non-adsorbable Ni(I)-cyclam - CO with accumulated CO in the vicinity of the growing electrode surface. At more negative potentials, the Hg surface became naked almost from the beginning of drop growth and inactive for CO, reduction. Further negative potentials evolved hydrogen just as in the absence of the catalyst. In-situ confirmation of the catalytic generation of CO at potentials between - 1.15 and - 1.70 V and hydrogen evolution at the more negative potentials during CV was successfully done in this laboratory [35] by differential electrochemical mass spectroscopy (DEMS) [36,37] with a modification - the use of an amalgamated gold mesh electrode. Electrocatalytie behaviour of Ni(l)-@am,,
on Hg-Au for CO, reduction
In the following, the catalytic reduction of CO2 on Hg-Au in the presence of Ni(II)-cyclam will be discussed on the basis of the experimental findings described so far.
212 0
5 t 2 7J j 0.5 E 3 ” Z ”
1.0
-2.0 E
-1.5 /
v
vs.
-1.0 SCE
- 0.5
Fig. 10. The calculated catalytic cu~nt-potential curves based on the EC mechanism under (a) N2, (b) low2 atm of CO2 and (c) 10-l atm of COz, i.e. potential dependence of the surface concentration of Ni(I)-cyclam, as a function of the electrode potential at various CO parfial pressures. The current or the surface concentration is normalized by that for fully reduced Ni(I)-cycfam,, under N,.
If we assume (i) no catalytic activity of the Ni(I)-cyclam in the bulk, (ii) a linear dependence of the catalytic activity on the surface concentration of the Ni(I)cyclam,,, (iii) independence of the catalytic activity of Ni(I)-cyclam, on the potential and (iv) a great excess of CO2 in the vicinity of the electrode surface with negligible depletion of CO, by the reduction, the I-E curve for CO, reduction will change as shown by curve a in Fig. 10. In the calculation, the surface concentration of Ni(I)-cyclist is further assumed to be independent of the electrode potential when the potentials are negative enough for all the adsorbed complexes to be reduced. The final assumption was made for brevity against the experimental result. Namely, the decrease in the difference in drop time at potentials more negative than - 1.56 V under N, in Fig. 5 apparently indicates a decrease in the surface concentration of Ni( I)-cyclam ad with the decrease in the potential. The shape of curve a should be the same as the charge (which was obtained by integrating the surface wave in Fig. lc) vs. potential curve, because the catalytic current is assumed to be proportional to the amount of Ni(I)-cyclam,. The first assumption as to negligible catalytic activity of Ni(I)-cyclam in solution in comparison with Ni(I)-cyclam,, was corroborated by the abrupt drop in the catalytic currents both in CV and in polarography under CO, when Ni(I)-cyclam,, desorbs at potentials more negative than - 1.70 V, as indicated by the drop time under C02. The polarographic outline under CO2 in Fig. 6c in the potential region between -1.15 and - 1.40 V is in rough agreement with the shape of curve a in Fig. 10. Therefore, the sigmoidal I-E curve in the polarogram reflects the increase in by the potential sweep in the adsorption surface concentration of Ni(I)-cyclam, prewave region. However, the further increase in current in CV up to - 1.50 V in Fig. 3d and the potential-dependent increase in the initial instantaneous currents up
213
to - 1.70 V in the I-t curves in Fig. 9 both give a good indication that the catalytic current increased not only with the Ni(I)-cyclam., concentration, but also with the applied further negative potential against the third assumption. When we take account of the potential-dependent decrease in adsorbability of Ni(I)-cyclam under N, in Fig. 5, the increase in the catalytic activity, even after full reduction of all the surface catalyst to the Ni(1) state, is clearly due to the increased negative potentials. The potential dependence of the catalytic activity of Ni(I)-cyclam., described above rules out a purely surface EC catalytic mechanism [38,39] for CO, reduction on this surface. Here, the purely EC catalytic mechanism means that the catalytic reaction rate depends only on the surface concentration of the electrochemically generated active catalyst (in the present case Ni(I)-cyclam,,) and is not affected by the potential, i.e. the I-E curve will follow curve a in Fig. 10 under the condition of excess CO,. It cannot be concluded, however, from the above-described results alone what kind of mechanism in terms of successive elementary reactions is applicable to the present CO2 reduction. Further detailed study will be needed for this point. Finally, we will discuss the question of why the amount of Ni(I)-cyclam., decreased when the reduction of Ni(II)-cyclam in the bulk started. In the presence of a large amount of electrochemically produced CO, which is the case for the CV in Fig. 3, Ni(I)-cyclam in the vicinity of the electrode surface reacts with CO rapidly to form Ni(I)-cyclam . CO. It is unlikely, however, that the carbonyl complex is an intermediate in the main route for CO formation, as proposed by Sauvage and co-workers [20], because the carbonyl complex is fairly stable in the time scale of CV with high sweep rates of 300-500 mV s-’ under CO. It is also very reasonable to assume that the carbonyl complex is inactive to CO,, because the reversible wave for the Ni(II)/Ni(I)-TMC redox couple under N, shifted towards positive potentials under CO, but did not shift under CO, [40], i.e. the coordination power of CO, to Ni(1) is much lower than that of CO. If the rapid formation, the stability and the inertness to CO2 of Ni(I)-cyclam . CO described above all hold in the time scale of the CV, the decrease in the catalytic current beyond the peak potentials with decreasing potential in Fig. 3d can be rationalized as follows. (i) During the negative potential sweep across the potential region of the adsorption prewave on Hg-Au, CO is produced catalytically and accumulated gradually in the vicinity of the electrode surface, probably according to the rate law in d[CO]*/dt
= k,,,[CO,]*
- [Ni(I)-cyclam,,]
(7) where the asterisk denotes concentrations in the vicinity of the electrode surface. (ii) When the potential becomes negative enough to reduce Ni(II)-cyclam in solution, the reduced species in solution reacts rapidly with electrogenerated CO according to eqn. (4). Hence, the concentration of the uncarbonylated surface species decreases through the equilibrium in reactions in the following equations: Ni(II)-cyclam Ni(II)-cyclam,,
+ Ni(II)-cyclam., + e-+
Ni(I)-cyclam,
(8) (9)
214
where the concentration of uncarbonylated Ni(II)-cyclam in the vicinity of the electrode surface may decrease with the increase in accumulated CO concentration according to the Nernst equation for reaction (4) as follows: E=
E” + (RT/nF)
ln{[Ni(II)-cyclam]*[CO]*/[Ni(I)-cyclam.CO]*}
(IO)
where E o is -1.39 V. (iii) If the concentration of electrogenerated CO is much higher than those of the Ni-cyclams, which is the case for CO* reduction by Ni(I)-cyclam., on Hg-Au, and is taken as approximately constant, the concentration of Ni(II)-cyclam can be calculated readily by eqn. (10). Then we can estimate the change in the concentration of Ni(I)-cyclam., from the adsorption isotherm, i.e. the charge for the adsorption prewave vs. the concentration of Ni(II)-cyclam, and finally the catalytic current, with the aid of eqn. (7) as a function of the applied electrode potential. The change in the catalytic current as a function of the concentration of Ni(II)-cyclam in the vicinity of the electrode surface, however, can be obtained directly by the curve in Fig. 4 to a fairly good approximation. In other words, the curve in Fig. 4 can be regarded as the adsorption isotherm by the relation in eqn. (7). In the above discussion, no carbonylation of Ni(I)-cyclam,,, which seems to be very important for the catalytic reduction of CO,, was assumed implicitly because of the coordination by mercury as a fifth ligand. This weak interaction of CO to the adsorbed Ni(1) species for Ni(II)-TMC was confirmed by the experimental finding under CO that almost no positive potential shift was observed for the adsorption peak compared with the large positive shift for the Ni(II)/Ni(I)-TMC redox couple in solution. The spiky adsorption prewave for Ni(II)-TMC enabled us to confirm its negligible potential shift by CO in spite of the overlap with the shifted main wave, but confirmation was difficult for Ni(II)-cyclam with a broad prewave. Curves b and c in Fig. 10 show the theoretical current-potential curves for the cases where the concentrations of catalytically generated CO are assumed to be one-hundredth and one-tenth of that under CO at atmospheric pressure, respectively. In these calculations, the potential dependence of k,,, was neglected for brevity. The curve in Fig. 4 was used as the adsorption isotherm for calculation of the decrease in the catalytic current by the diminished concentration of Ni(II)-cyclam in the vicinity of the electrode surface. The qualitative resemblance between the experimental I-E curve and the calculated one indicates the validity of the interpretation of the decreased catalytic activity described above.
ACKNOWLEDGEMENTS
We would like to thank Dr. U. Akiba for helpful discussions and Mr. Y. Nakamura for assistance in part of the experimental work. This work was partially supported by a Grant-in-Aid for Scientific Research on Priority Area 63604534 and 63612506 from the Ministry of Education, Science and Culture.
215 REFERENCES 1 I. Taniguchi in J.O’M. Bockris (Ed.), Modem Aspects of Electrochemistry, No. 20: Electrochemical and Photoelectrochemical Reduction of Carbon Dioxide, Plenum Press, New York, 1989, p. 327. 2 J.P. ColIin and J.P. Sauvage, Coord. Chem. Rev., 93 (1989) 245. 3 S. Meshitsuka, M. Ichikawa and K. Tamaru, Chem. Commun., (1974) 158. 4 K. Hiratsuka, K. Takahashi, H. Sasaki and S. Toshima, Chem. Lett., (1977) 1137. 5 K. Takahashi, K. Hiratsuka, H. Sasaki and S. Toshima, Chem. Lett., (1979) 305. 6 B. Fisher and R. Eisenberg, J. Am. Chem. Sot., 102 (1980) 7361. 7 P.A. Christensen, A. Hamnett and A.V.G. Muir, J. Electroanal. Chem., 241 (1988) 361. 8 N. Furuya and K. Matsui, J. Electroanal. Chem., 271 (1989) 181. 9 S. Kapusuta and N. Hackerman, J. Electrochem. Sot., 131 (1984) 1511. 10 M. Hannnouche, D. Lexa and J.M. Savtant, J. Electroanal. Chem., 249 (1988) 347. 11 C.M. Lieber and N.S. Lewis, J. Am. Chem. Sot., 106 (1984) 5033. 12 M.G. Bradley, T. Tysak, D.J. Graves and N.A. Vlachopoulos, Chem. Cornmun., (1983) 349. 13 A.H. Tinnemans, T.P.M. Koster, D.H.M.W. Thewissen and A.M. Mackor, Reel. Trav. Chim. Pays-Bas, 103 (1984) 288. 14 M. Tezuka, T. Yajima, A. Tsuchiya, Y. Matsumoto, Y. Uchida and M. Hidai, J. Am. Chem. Sot., 104 (1982) 6834. 15 J. Hawecker, J.-M. Lehn and R. Ziessel, Chem. Commun., (1984) 328. 16 S. Slater and J.H. Wagenknecht, J. Am. Chem. Sot., 106 (1984) 5367. 17 C.M. Bohnger, B.P. Sullivan, D. Conrad, J.A. Gilbert, N. Story and T.J. Meyer, Chem. Commun., (1985) 796. 18 S. Daniele, P. Ugo, G. Bontempelli and M. Fiorani, J. Electroanal. Chem., 219 (1987) 259. 19 M. Beley, J.-P. Collin, R. Ruppert and J.-P. Sauvage, Chem. Commun., (1984) 1315. 20 M. Beley, J.-P. Collin, R. Ruppert and J.-P. Sauvage, J. Am. Chem. Sot., 108 (1986) 7461. 21 M. Beley, J.-P. Collin, J.-P. Sauvage, J.-P. Petit and P. Chartier, J. Electroanal. Chem., 206 (1986) 333. 22 J.-P. Petit, P. Chartier, M. Beley and J.-P. Deville, J. Electroanal. Chem., 269 (1989) 267. 23 J.L. Grant, K. Goswami, L.O. Spreer, J.W. Otvos and M. Calvin, J. Chem. Sot., Dalton Trans., (1987) 2105. 24 A.J. Bard and L.R. Faulkner, Electrochemical Methods, Wiley, New York, 1980. 25 J. Heyrovsky and J. Kuta, Principles of Polarography, Academic Press, New York, 1966. 26 M. Ciampolini, L. Fabbrizzi, M. Licchelli, A. Perotti, F. Pezzini and A. Poggi. Inorg. Chem., 25 (1986) 4131. 27 M. Ciampolini, L. Fabbrizzi, A. Perotti, A. PO&, B. Seghi and F. Zanobini, Inorg. Chem., 26 (1987) 3527. 28 K. Barefield and F. Wagner, Inorg. Chem., 12 (1973) 2435. 29 J.-P. Collin, A. Jouaiti and J.-P. Sauvage, Inorg. Chem., 27 (1988) 1986. 30 N. Jubran, G. G&burg, H. Cohen, Y. Koresh and D. Meyerstein, Inorg. Chem., 24 (1985) 251. 31 C.M. Madeyski, J.P. Michael and R.D. Hancock, Inorg. Chem., 23 (1984) 1487. 32 Y. Hirata, Y. Nakamura, U. Akiba, K. Suga and M. Fujihira, submitted. 33 R.R. Gagnt and D.M. Ingle, Inorg. Chem., 20 (1981) 420. 34 M. Kodama and E. Kimura, J. Chem. Sot., Dalton Trans., (1976) 2335. 35 H. Hirata, K. Suga and M. Fujibira, Chem. Lett., (1990) 1155. 36 0. Wolter and J. Heitbaum, Ber. Bunsenges. Phys. Chem., 88 (1984) 2. 37 D. Tegtmeyer, A. Heindrichs and J. Heitbaum, Ber. Bunsenges. Phys. Chem., 93 (1989) 201. 38 J. Zak and T. Kuwana, J. Electroanal. Chem., 150 (1983) 645. 39 M. Fujihira in A.J. Fry and W.E. Britton (Eds.), Topics in Organic Electrochemistry: Modified Electrodes, Plenum Press, New York, 1986, p. 255. 40 Y. Hirata, K. Suga and M. Fujihira, Thin Solid Films, 179 (1989) 95.
J. Electroanal. Chem., 292 (1990) 199-215 Elsevier Sequoia S.A., Lausanne
Electrocatalytic
reduction of CO, by nickel(I1) cyclam
Study of the reduction mechanism on mercury by cyclic voltammetry, polarography and electrocapillarity Masamichi Fujihira *, Yoshiki Hirata and Kosaku Suga Department of Biomolecular Yokohama 227 (Japan) (Received
29 March
Engineering,
Tokyo Institute
of Technology,
4259 Nagatsuta,
Midori-ku,
1990; in revised form 9 May 1990)
ABSTRACT The electrocatalytic reduction of CO1 on mercury with Ni(II)-cyclam was studied in detail by cyclic voltammetry, polarography and electrocapillarity of aqueous solutions with and without the catalyst under N,, CO and COa. It is concluded from these electrochemical measurements that (i) the adsorbed complex (Ni(I)-cyclam,,) on Hg, not the reduced complex in solution Ni(I)-cyclam, is the active catalyst for the CO* reduction; (ii) both Ni(II)- and Ni(I)-cyclam can adsorb on Hg, and the oxidation states of the adsorbed complexes depend on the electrode potentials; (iii) Ni(I)-cyclam adsorbs on Hg even at - 2.0 V (vs. SCE) under N,, but the surface concentration of Ni(I)-cyclam,, decreases gradually with a decrease in the potential and desorbs completely at potentials more negative than - 1.7 V under CO or under CO or CO, is due to the formation of COa; and (iv) the desorption of Ni(I)-cyclam,, unadsorbable Ni(I)-cyclam.CO by the reaction of Ni(I)-cyclam in solution with dissolved CO or electrocatalytically generated CO from COa.
INTRODUCTION
The electrochemical reduction of CO, on metallic cathodes requires typically negative potentials more than- -2.0 V vs. SHE (- 2.24 V vs. SCE) [l]. Energy consumption is reduced by the use of molecular electrocatalysts [2] such as metallophthalocyanines [3-S], metalloporphyrins [9,10] and various tetraazamacrocyclic complexes [ll]. CO or formate ion was shown to be the major product of the reduction [5,6,11]. Several macrocyclic complexes have also been used for the photoelectochemical reduction of CO, to CO with various semiconductors at rela-
l
To whom correspondence
0022-0728/90/$03.50
should
be addressed.
0 1990 - Elsevier Sequoia
S.A.
200
tively low negative potentials [12,13]. Various acyclic complexes have also been proposed as electrocatalysts, including an iron-sulphur cluster [14], and complexes of rhenium [15], rhodium [16,17], ruthenium [17] and nickel [18]. Sauvage and co-workers have reported the exceptional efficiency and selectivity of nickel(I1) cyclam (cyclam is an abbreviation for 1,4,8,11-tetraazacyclotetradecane) as an electrocatalyst for the reduction of CO, on Hg in water [19,20]. The catalytic reduction involves the production of CO by the following electrochemical reaction [1,2]: CO,+2H++2e-+CO+H,O
E”‘=
-0.76V(vs.SCE)
(I)
This catalyst has also been used for the photoelectrochemical [21,22] and photochemical [23] reduction of CO,. However, the details of the electrochemical reduction mechanism of CO, by Ni(II)-cyclam have not yet been clarified. In particular, present knowledge of catalytically active species that are adsorbed on Hg is limited. Therefore, in this work the role of the adsorbed catalyst on the electrochemical reduction of CO, on Hg was studied in detail by cyclic voltammetry (CV) [24], polarography [25] and electrocapillarity [25] in an aqueous Ni(II)-cyclam solution under various atmospheric conditions (N,, CO and CO,). As a reference, the cyclic voltammetric behaviour on glassy carbon (GC) was also studied. EXPERIMENTAL
Cyclam was received from Aldrich. N-Tetramethylated cyclam (1,4,8,11_tetramethyl-1,4,8,11-tetraazacyclotetradecane, abbreviated as TMC) was prepared as described by Ciampolini et al. [26,27]. The nickel(I1) complexes with these ligands were prepared and obtained as perchlorate according to the literature [28]. The following gases (Tomoe) were used without further purification: nitrogen (99.9995%), carbon monoxide (99.9%) and carbon dioxide (99.99%). Aqueous solutions were prepared from analytical grade reagents and pure water from a Milli-Q system (Millipore Ltd.). The pH of the solution under CO, was adjusted by adding NaHCO, to aqueous 0.1 M NaClO,. Under N, or CO, an aqueous solution of 0.1 M NaClO, without buffer was used as the electrolyte. CV, polarography and electrocapillarity were carried out with a Nikko Keisoku potentiostat by using an electrochemical cell with a salt bridge containing aqueous 0.1 M NaClO, in order to remove chloride contamination from the saturated calomel electrode (SCE) used as the reference electrode. An amalgamated gold (Hg-Au) disk (1.5 mm in diameter), a CC disk (3.0 mm in diameter) or a dropping mercury electrode (DME) was used as the working electrode together with a GC rod counter-electrode. The Hg-Au disk electrode was prepared by polishing a Au disk electrode with emery paper (No. 2000) and then 0.3 pm alumina, sonicating it with Milli-Q water and finally amalgamating it by dipping the electrode into GR-grade Hg. The GC disk electrode was also polished and sonicated similarly prior to use. All measurements were done at room temperature (ca. 20°C). All the potentials given below are referred to the SCE.
201 RESULTS
AND
DISCUSSION
Cyclic voltammetry
of Ni(II)-cyclam
itself in water under N, on GC and Hg-Au
The cyclic voltammograms of Ni(II)-cyclam in an aqueous solution containing 0.1 M NaClO, under N, on the GC and Hg-Au disk electrodes are shown in Figs. la and lb, respectively. Two reversible waves at + 0.80 and - 1.56 V, corresponding to the Ni(III)/Ni(II)and Ni(II)/Ni(I)-cyclam redox couples respectively, were observed on GC; while only one reversible wave at - 1.56 V for the Ni(II)/Ni(I)cyclam couple could be recorded on Hg-Au because of Hg dissolution at positive potentials. The sweep rate (v) dependence of the peak heights of these reversible waves, i.e. a linear dependence of the peak currents on z?/*, indicates clearly their diffusion-controlled characteristics [24]. The height of the reoxidation peak of the Ni(II)/Ni(I)-cyclam redox couple was slightly decreased at slow sweep rates,
10pA
(b)
(cl
0.8jJA I
-2.0
-1.0 E
0 /
v
vs.
1.0 SCE
Fig. 1. Cyclic voltammograms of 1 mM Ni(II)-cyclam aqueous solution containing 0.1 M NaCIO, under N, on (a) a GC disk electrode and (b) a Hg-Au disk electrode. (c) Sweep rate dependence of the adsorption prewave of the same solution on the same Hg-Au disk electrode. Sweep rate/V s-l: (a, b) 0.1; (c) 0.1, 0.2, 0.3, 0.4, 0.5.
202
probably [29]:
because of the catalytic
2 Ni(I)-cyclam++
nature of Ni(I)-cyclam
2 H,O + 2 Ni(II)-cyclam2++
and/or the limited lifetime of Ni(I)-cyclam Ni(I)-cyclam++
for hydrogen evolution
H, + 2 OH-
(2)
due to a ligand loss reaction [30]:
2 H,O + Ni(I),‘, + cyclamH:+
+ 2 OH-
(3)
The reduced overpotential for hydrogen evolution upon repeated potential cycling was more apparent on GC than on Hg-Au. The deposition of nickel on GC by the electrochemical reduction of Ni(I)& generated in eqn. (3) may account for the activation of hydrogen evolution. It is also interesting to note that one small reversible prewave of the Ni(II)/Ni(I)-cyclam redox couple described previously by Sauvage and co-workers [20] was observed clearly only on Hg-Au. Figure lc shows the sweep rate dependence of the prewave observed at ca. - 1.33 V on Hg-Au in 1 mM Ni(II)-cyclam aqueous solution. The almost linear dependence of the peak height on the sweep rate indicates that this wave can be ascribed to the adsorbed Ni(II)/Ni(I)-cyclam (Ni(II)/Ni(I)-cyclam.,) redox couple on a Hg-Au electrode surface [24]. The positive shift of the adsorption peak potential compared with the half-wave potential of the redox couple in solution at -1.56 V shows clearly that Ni(I)-cyclam adsorbs more strongly on Hg than does Ni(II)-cyclam. The importance of Ni(I)cyclam., in the catalytic reduction of CO, will be described later. Cyclic voltammetry
of Ni(II)-cyclam
itself in water under CO on GC and Hg-Au
In order to clarify the electrocatalytic reduction mechanism of CO,, the cyclic voltammetric behaviour of Ni(II)-cyclam under CO was examined prior to CV under CO,. CO was found to be the main product of CO, electrolyses on Hg in the presence of Ni(II)-cyclam in water [20]. In Fig. 2a a cyclic voltammogram of 1 mM Ni(II)-cyclam in 0.1 M NaClO, under CO on GC between - 1.50 and + 1.15 V is shown. The potential sweep was started from - 1.0 V in the positive direction. During the first cycle (solid line) between -1.0 and +1.15 V, no change was observed between the cyclic voltammograms under N, and under CO. Namely, only one reversible wave was observed at the same potential of + 0.80 V as that observed under N, which was responsible for the Ni(III)/Ni(II)-cyclam redox couple without any further coordination reaction by CO. In contrast, the other reversible wave corresponding to the Ni(II)/Ni(I)-cyclam redox couple under N, was shifted under CO in the positive direction by 170 mV and its reoxidation peak height was further diminished. In the second positive sweep (dashed line), two new small peaks appeared at ca. - 0.6 and + 1.0 V. When the lower end of the potential sweep was extended to - 1.8 V as shown in Fig. 2b, another new anodic wave appeared at +0.6 V and its peak height increased with repeated potential cycling. The origin of the new peaks observed at - 0.6, + 0.6 and +l.O V has not yet been clarified, but Sauvage and co-workers [20] found that the
203
Fig. 2. Cyclic voltammograms of 1 mM Ni(II)-cyclam aqueous solution containing 0.1 M NaCIO, under CO on a CC disk electrode with (a) a narrow potential sweep and (b) a wide potential sweep, and (c) on a Hg-Au disk electrode with a narrow potential sweep. Sweep rate: 0.1 V s-l. The solid lines and dashed tines indicate the first and second cycle, respectively.
Ni(1) carbonyl complex described below is much less stable in H,O than in DMF. By taking account of a similar positive shift by CO observed for the more stable redox couple of Ni(II)/Ni(I)-TMC [30-321, the following coordination reaction of CO to electroproduced Ni(I)-cyclam is most likely responsible for the positive potential shift: Ni(II)-cyclam++
e-+
CO + Ni(I)-cyclam.
CO
(4)
Several tetraazamacrocyclic complexes of Ni(1) having carbonyl as the fifth ligand with square-pyramidal geometry have been generated in solution [33]. As no new peaks appeared for the CV of Ni(II)-TMC in the presence of CO, except for the potential shift of the reversible wave at -0.99 V under N, to -0.75 V under CO [32], the carbonyl formation in eqn. (4) cannot be responsible for the new anodic peaks. The same potential shift as for the Ni(II)/Ni(I)-cyclam redox couple under CO was also observed on Hg-Au as shown in Fig. 2c. When the sweep rate was changed from 100 to 500 mV s-l, there was almost no change in the shifted half-wave potential at -1.39 V or in the peak separation between the cathodic and anodic
204
peaks. Rather, an increase in the reoxidation wave-height of the redox wave accompanied by a relative decrease in the height of the anodic wave of a by-product at ca. -0.7 V was observed with the increase in sweep rate. These observations indicate that the coordination reaction in eqn. (4) is very rapid and fairly reversible. As is obvious from a comparison of Figs. 2a and 2c, the peak potential of the by-product was shifted from - 0.6 to - 0.7 V upon changing the electrode material from GC to Hg-Au. In addition to this anodic peak, another new sharp anodic peak at - 0.06 V and a new redox wave at + 0.05 V appeared. The peak currents of the latter waves were characteristic on Hg-Au and increased with repeated potential cycling or by extending the lower potential limit. As these waves were not observed on GC, a reaction between some other unknown by-products of Ni(II)-cyclam and Hg is most likely the cause of these new waves near 0 V on Hg-Au. CV of 1 mM free ligand cyclam on Hg-Au under N, gave a redox wave at - 0.07 V. This wave can be related to Hg(II)-cyclam/Hg’ redox couples [31,34]. Therefore, the redox waves observed at +0.05 V under CO in Fig. 2c suggest that, in the presence of CO, the ligand loss reaction in eqn. (3) is accelerated at negative potentials and the resulting free cyclam gives a redox wave for Hg(II)-cyclam/Hg’ which is shifted slightly to positive potentials by carbonyl formation.
Cyclic voltammetty
of Ni(II)-cyclam
in water under CO, on GC and Hg-Au
The cyclic voltammograms of 1 mM Ni(II)-cyclam under CO, in aqueous 0.1 M NaClO, containing 35 mM NaHCO, (pH 6.2) on GC and Hg-Au are shown in Figs. 3c and d, respectively. In Figs. 3a and 3b, their background currents observed in the same electrolyte under CO, without Ni(II)-cyclam (pH 6.2) are shown. On GC, two cathodic peaks are seen at - 1.5 and - 1.8 V in the negative sweep, while one cathodic peak at - 1.8 V and one small anddic peak at - 0.7 V are seen in the positive sweep. On Hg-Au, one large cathodic peak at -1.5 V in the negative sweep and two small anodic peaks at ca. -0.7 and -0.1 V in the positive sweep are seen. In comparison with their background currents and the cyclic voltammograms of Ni(II)-cyclam under N, and CO, the catalytic reduction of CO, in the presence of Ni(II)-cyclam is obvious. But the efficiency on GC was found to be much lower than that on Hg-Au in terms of the current density in A cme2, i.e. if we take account of the four times larger area of GC than that of Hg-Au. In addition to this magnitude of the catalytic current density, the onset potential for the catalytic current was much more favourable on Hg-Au than on GC. This effect of the substrate electrode material on the catalytic nature supports strongly the view that the adsorbed species plays an important role in the catalysis. The appearance of the anodic peaks under CO, in Figs. 3c and 3d are ascribable to the by-products, which were not observed under N, (Fig. 1) but were observed under CO (Fig. 2) and suggests that the main reduction product of CO, is CO on both electrode materials.
205
I
I
,
I
I
-2.0 E
/
I
-1.0 V
I
a
I
8
I
t
0 vs.
SCE
Fig. 3. Cyclic voltammograms of aqueous solutions containing 0.1 M NaClO, and 0.035 M NaHCO, (pH 6.2) under CO2 on a GC disk electrode (a) without and (c) with 1 mM Ni(II)-cyclam, and on a Hg-Au disk electrode (b) without and (d) with 1 mM Ni(II)-cyclam. Sweep rate: 0.1 V s-l.
Dependence
of the catalytic
currents on the concentration
of Ni(II)-cyclam
in solution
As is apparent from the comparison between the adsorption prewave of Ni(II)cyclam in Fig. lc and the catalytic reduction wave of CO, in Fig. 3d, the two onset potentials for the adsorption and the catalytic waves are in good agreement. This coincidence implies that Ni(I)-cyclam, on Hg is responsible for the catalysis, as Sauvage and co-workers first pointed out [20]. We also repeated CV on Hg-Au under CO, in aqueous solutions containing various concentrations of Ni(II)-cyclam. As shown in Fig. 4, we could reproduce an adsorption isotherm-like dependence of the peak intensity on the logarithmic concentration of Ni(II)-cyclam in M. Again, this observation supports strongly the importance of Ni(I)-cyclam., for the catalytic reduction of CO, in comparison with Ni(I)-cyclam formed in the aqueous bulk.
206
log [ NiW-cyclam
I
Fig. 4. Plot of the cathodic peak current density of CV for the catalytic electrode vs. the logarithmic concentration of Ni(II)-cyclam in M.
CO, reduction
on a Hg-Au
disk
From the cyclic voltammetric behaviour of Ni(II)-cyclam on Hg under N,, CO and CO, described so far, it is not clear whether Ni(II)-cyclam was adsorbed on Hg as well as Ni(I)-cyclam. The potential dependence of the amount of Ni(I)-cyclam., is also not obvious from CV. A study of the electrocapillarity was attempted, therefore, for Hg/aqueous 0.1 M NaClO, without Ni(II)-cyclam and Hgfaqueous 1 mM Ni(II)-cycIam in the supporting electrolyte systems under N,, CO and CO,. The electrocapillary curves were obtained approximately by measuring the average drop time for ten successive drops, keeping the applied electrode potential and the height of Hg column constant, and by repeating the drop time measurements at different potentials. The typical drop time curves thus obtained are shown in Fig. 5. In the absence of Ni(II)-cyclam, an approximately parabolic curve was obtained under CO which coincided with the curves under N, and COz, indicating no appreciable adsorption of CO on Hg. In contrast, in the presence of Ni(II)-cyclam a decrease in the drop time from that observed without the catalyst was seen under N, at potentials more negative than -0.3 V. The difference in drop time increased gradually with decreasing electrode potential until the potential reached the redox potential for Ni(II)/Ni(I)cyclam at - 1.56 V. Beyond this potential, the difference in drop time decreased with a decrease in the potential, but it was still lower than that observed without the catalyst even at - 2.0 V, It is interesting to note that a break in the curve was not
201
vv V
(a)
0
(b) (cl (dl
A 0 V
V V V %
V 0 %7 0
IQ
40
0
V V
0 “Ooo 0 3.0
%
Q
,,,,,,,,,,,,,,,,,,,P~ 0
-0.5 E
-1.5
-1.0 /
V
vs.
-2.0
SCE
Fig. 5. Drop time curves of aqueous 0.1 M NaCIO, solutions with 1 mM Ni(II)-cyclam under (a) N,, (b) CO and (c) CO2 with 0.035 M NaHCO, (pH 6.2), and (d) without Ni(II)-cyclam and 0.035 M NaHC4 under N,.
observed at the potential for Ni(II)/Ni(I)-cyclam.,, but at the potential for Ni(II)/Ni(I)-cyclam in solution. The decrease in drop time observed at all potentials more negative than - 0.3 V indicates clearly that both Ni(II)- and Ni(I)-cyclams can adsorb on Hg, although the redox states of the adsorbed complexes, of course, depend on the electrode potential. Under CO, a difference in the drop time with and without the catalyst was also observed at potentials less than -0.3 V, but a break in the curve was found at ca. -1.35 V which corresponded to the half-wave potential for the reduction of Ni(II)-cyclam in the presence of CO, i.e. the reaction described in eqn. (4). At potentials of ca. - 1.7 V or less, almost no difference in drop time with or without the catalyst was observed. The difference between the drop time curves under N, and CO implies that coordination of CO to the electrogenerated Ni(I)-cyclam in solution diminishes the amount of Ni(I)-cyclamad. Under CO,, the drop time curve was very similar to that observed under CO. The lack of adsorbability of Ni(I)-cyclam . CO produced by the reaction between Ni(I)cyclam in solution and the electrocatalytically generated CO may also account for the curve under CO,. Polarographic
study of the catalytic reduction of CO, on mercury
In Figs. 6a-6c are shown the polarograms in instantaneous currents of 1 mM Ni(II)-cyclam under N,, CO and CO,, respectively. Under N,, a reversible wave
208
(a)
fb)
III1
I
I
-2.0
I
Iii!
i
-1.5 E
Fig. 6. Polarograms
/
I
I
- 1.0 v
vs.
SCE
of 1 mM Ni(II)-cyclam
CO and (c) CO2 with 0.035 M NaHCO,
aqueous
solution
containing
0.1 M
NaClO.,under(a) NZ,(b)
(pH 6.2).
with a half-wave potential of -1.56 V can be seen, but an adsorption prewave, which is well known for methylene blue [25], is not obvious in Fig. 6a. The adsorption prewave for Ni(II)-cyclam could be barely seen under high magnification of the current axis. The reversible wave was shifted to a positive potential by carbonyl formation in the same way as in CV under CO (Fig. 6b). However, the polarographic wave of the catalytic reduction of CO, was complex, as in CV. Maximum currents were not necessarily observed at the end of growth of the DME. The outline connecting the maximum currents is similar to the cyclic voltammogram in Fig. 3d; that is, the decrease in the currents on the negative side of the peak potential of CO, in CV was also observed in Fig. 6c. This immediately suggests that the abruptly decreased current in CV in Fig. 3d was not due to the depletion of CO,, but due mainly to the decrease in catalytic activity of the Hg surface with Ni(I)-cyclam,. An excess of C&, however, is not enough for the rate to be regarded as a quasi-first-order reaction, SO that the
209 b +
a
d +
L !L jL f +
+
L
JL
5
JL j + ?f
;
: -L/J
I
~OJJA
h
‘2;. Fig. 7. Changes in the instantaneous current as a function polarography of 1 mM Ni(II)-cyclam aqueous solution containing under CO1 (pH 6.2). Applied electrode potential/V (vs. SCE): -1.40, (e) -1.45, (f) -1.50, (g) -1.55, (h) -1.60, (i) -1.65, - 1.85, (n) - 1.90, (0) - 1.95.
of time at various potentials in the 0.1 M NaClO, and 0.035 M NaHCO, (a) -1.25, (b) -1.30, (c) -1.35, (d) (j) -1.70, (k) -1.75, (1) -1.80, (m)
catalytic peak current at ca. - 1.4 V decreased slightly with the decrease in height of the Hg column indicating a small effect of depletion of CO*. It is worth noting that the potential ca. - 1.7 V where a sudden decrease in current was observed in the polarogram is in good agreement with the potential where adsorption of the catalyst ended under CO, as indicated by the drop time curve in Fig. 5. To study the polarogram in detail, the dependence of the instantaneous current on time, i.e. I-t curves, was observed as a function of the applied electrode potential. Typical I-t curves recorded from dropping of the preceding DME with a pen recorder are shown in Fig. 7. At the potentials between the onset of ca. - 1.15 V and - 1.40 V giving the maximum catalytic current, the instantaneous current (1) increased smoothly with time (t). Log I vs. log t plots in this potential region are shown in Fig. 8. At all potentials, the slopes were initially 2/3, which corresponds to the increase in the surface area of the DME with time [25]: A = 0 85m2/3t2/3
(5)
where m is the flow rate of mercury in g s-l. But the slopes can be seen to deviate from this theoretical slope afterwards, during the drop-life. This deviation started at an earlier stage of drop growth and became larger with an increase in the applied negative potential. The proportionality of the catalytic current to the surface area
210
Fig. 8. Log I-log t plots for changes in instantaneous currents as a function of time at various potentials in V vs. SCE: (a) -1.40, (b) - 1.35, (c) -1.30, (d) -1.25.
gives a good indication of the rapid adsorption equilibrium of the catalyst and the surface catalytic nature of the CO, reduction. Depletion of the CO, concentration in the vicinity of the electrode surface is most likely the cause of the deviation from the theoretical slope in this potential region. As is obvious from the I-t curves at further negative potentials in Fig. 7, the rise in the initial current became increasingly steeper, but the current levelled off and even fell in the last seconds of the drop-life. At - 1.70 V, a rapid rise and fall in I was observed in the initial 1 s. At further negative potentials, the currents were kept low during their drop-lives. At potentials more negative than -1.80 V, again a monotonic increase in the current with time was observed. The current observed in this most negative potential region is attributable to hydrogen evolution, described later. In Fig. 9, the instantaneous currents at 0.3, 0.5 and 3.0 s after the beginning of drop growth are plotted as a function of the applied electrode potential. From the potential dependence of the current observed at 0.3 s, the catalytic reduction of CO, was found to be accelerated by the increase in negative potential until the potential reached - 1.7 V, where the catalyst was desorbed in the presence of catalytically generated CO. If the potential were negative enough for the observed current to be
211
-200 -2R
E
/
-1.5 V vs.
-1.0 SCE
Fig. 9. Instantaneouscurrentsat (a) 0.3 s, (b) 0.5 s and (~1 3.0 s after the beginningof drop growth as a function of the electrode potentialplotted from the 1-f curves in Fig. 7.
diffusion controlled of CO,, the I-t curve would behave as predicted by the IlkoviE equation [25]: 1, = 0.732nF~~‘/~~~/~~‘/~
(6)
and the current for each growing drop would increase with time during the drop-life. Therefore the decrease in the instantaneous current with time observed at potentials between - 1.40 and - 1.70 V should be attributed not to CO, depletion, but to the gradual decrease in the surface catalyst concentration during the drop-life. The decrease with time in the amount of Ni(I)-cyclam=~ may be due to the increase in the concentration of non-adsorbable Ni(I)-cyclam - CO with accumulated CO in the vicinity of the growing electrode surface. At more negative potentials, the Hg surface became naked almost from the beginning of drop growth and inactive for CO, reduction. Further negative potentials evolved hydrogen just as in the absence of the catalyst. In-situ confirmation of the catalytic generation of CO at potentials between - 1.15 and - 1.70 V and hydrogen evolution at the more negative potentials during CV was successfully done in this laboratory [35] by differential electrochemical mass spectroscopy (DEMS) [36,37] with a modification - the use of an amalgamated gold mesh electrode. Electrocatalytie behaviour of Ni(l)-@am,,
on Hg-Au for CO, reduction
In the following, the catalytic reduction of CO2 on Hg-Au in the presence of Ni(II)-cyclam will be discussed on the basis of the experimental findings described so far.
212 0
5 t 2 7J j 0.5 E 3 ” Z ”
1.0
-2.0 E
-1.5 /
v
vs.
-1.0 SCE
- 0.5
Fig. 10. The calculated catalytic cu~nt-potential curves based on the EC mechanism under (a) N2, (b) low2 atm of CO2 and (c) 10-l atm of COz, i.e. potential dependence of the surface concentration of Ni(I)-cyclam, as a function of the electrode potential at various CO parfial pressures. The current or the surface concentration is normalized by that for fully reduced Ni(I)-cycfam,, under N,.
If we assume (i) no catalytic activity of the Ni(I)-cyclam in the bulk, (ii) a linear dependence of the catalytic activity on the surface concentration of the Ni(I)cyclam,,, (iii) independence of the catalytic activity of Ni(I)-cyclam, on the potential and (iv) a great excess of CO2 in the vicinity of the electrode surface with negligible depletion of CO, by the reduction, the I-E curve for CO, reduction will change as shown by curve a in Fig. 10. In the calculation, the surface concentration of Ni(I)-cyclist is further assumed to be independent of the electrode potential when the potentials are negative enough for all the adsorbed complexes to be reduced. The final assumption was made for brevity against the experimental result. Namely, the decrease in the difference in drop time at potentials more negative than - 1.56 V under N, in Fig. 5 apparently indicates a decrease in the surface concentration of Ni( I)-cyclam ad with the decrease in the potential. The shape of curve a should be the same as the charge (which was obtained by integrating the surface wave in Fig. lc) vs. potential curve, because the catalytic current is assumed to be proportional to the amount of Ni(I)-cyclam,. The first assumption as to negligible catalytic activity of Ni(I)-cyclam in solution in comparison with Ni(I)-cyclam,, was corroborated by the abrupt drop in the catalytic currents both in CV and in polarography under CO, when Ni(I)-cyclam,, desorbs at potentials more negative than - 1.70 V, as indicated by the drop time under C02. The polarographic outline under CO2 in Fig. 6c in the potential region between -1.15 and - 1.40 V is in rough agreement with the shape of curve a in Fig. 10. Therefore, the sigmoidal I-E curve in the polarogram reflects the increase in by the potential sweep in the adsorption surface concentration of Ni(I)-cyclam, prewave region. However, the further increase in current in CV up to - 1.50 V in Fig. 3d and the potential-dependent increase in the initial instantaneous currents up
213
to - 1.70 V in the I-t curves in Fig. 9 both give a good indication that the catalytic current increased not only with the Ni(I)-cyclam., concentration, but also with the applied further negative potential against the third assumption. When we take account of the potential-dependent decrease in adsorbability of Ni(I)-cyclam under N, in Fig. 5, the increase in the catalytic activity, even after full reduction of all the surface catalyst to the Ni(1) state, is clearly due to the increased negative potentials. The potential dependence of the catalytic activity of Ni(I)-cyclam., described above rules out a purely surface EC catalytic mechanism [38,39] for CO, reduction on this surface. Here, the purely EC catalytic mechanism means that the catalytic reaction rate depends only on the surface concentration of the electrochemically generated active catalyst (in the present case Ni(I)-cyclam,,) and is not affected by the potential, i.e. the I-E curve will follow curve a in Fig. 10 under the condition of excess CO,. It cannot be concluded, however, from the above-described results alone what kind of mechanism in terms of successive elementary reactions is applicable to the present CO2 reduction. Further detailed study will be needed for this point. Finally, we will discuss the question of why the amount of Ni(I)-cyclam., decreased when the reduction of Ni(II)-cyclam in the bulk started. In the presence of a large amount of electrochemically produced CO, which is the case for the CV in Fig. 3, Ni(I)-cyclam in the vicinity of the electrode surface reacts with CO rapidly to form Ni(I)-cyclam . CO. It is unlikely, however, that the carbonyl complex is an intermediate in the main route for CO formation, as proposed by Sauvage and co-workers [20], because the carbonyl complex is fairly stable in the time scale of CV with high sweep rates of 300-500 mV s-’ under CO. It is also very reasonable to assume that the carbonyl complex is inactive to CO,, because the reversible wave for the Ni(II)/Ni(I)-TMC redox couple under N, shifted towards positive potentials under CO, but did not shift under CO, [40], i.e. the coordination power of CO, to Ni(1) is much lower than that of CO. If the rapid formation, the stability and the inertness to CO2 of Ni(I)-cyclam . CO described above all hold in the time scale of the CV, the decrease in the catalytic current beyond the peak potentials with decreasing potential in Fig. 3d can be rationalized as follows. (i) During the negative potential sweep across the potential region of the adsorption prewave on Hg-Au, CO is produced catalytically and accumulated gradually in the vicinity of the electrode surface, probably according to the rate law in d[CO]*/dt
= k,,,[CO,]*
- [Ni(I)-cyclam,,]
(7) where the asterisk denotes concentrations in the vicinity of the electrode surface. (ii) When the potential becomes negative enough to reduce Ni(II)-cyclam in solution, the reduced species in solution reacts rapidly with electrogenerated CO according to eqn. (4). Hence, the concentration of the uncarbonylated surface species decreases through the equilibrium in reactions in the following equations: Ni(II)-cyclam Ni(II)-cyclam,,
+ Ni(II)-cyclam., + e-+
Ni(I)-cyclam,
(8) (9)
214
where the concentration of uncarbonylated Ni(II)-cyclam in the vicinity of the electrode surface may decrease with the increase in accumulated CO concentration according to the Nernst equation for reaction (4) as follows: E=
E” + (RT/nF)
ln{[Ni(II)-cyclam]*[CO]*/[Ni(I)-cyclam.CO]*}
(IO)
where E o is -1.39 V. (iii) If the concentration of electrogenerated CO is much higher than those of the Ni-cyclams, which is the case for CO* reduction by Ni(I)-cyclam., on Hg-Au, and is taken as approximately constant, the concentration of Ni(II)-cyclam can be calculated readily by eqn. (10). Then we can estimate the change in the concentration of Ni(I)-cyclam., from the adsorption isotherm, i.e. the charge for the adsorption prewave vs. the concentration of Ni(II)-cyclam, and finally the catalytic current, with the aid of eqn. (7) as a function of the applied electrode potential. The change in the catalytic current as a function of the concentration of Ni(II)-cyclam in the vicinity of the electrode surface, however, can be obtained directly by the curve in Fig. 4 to a fairly good approximation. In other words, the curve in Fig. 4 can be regarded as the adsorption isotherm by the relation in eqn. (7). In the above discussion, no carbonylation of Ni(I)-cyclam,,, which seems to be very important for the catalytic reduction of CO,, was assumed implicitly because of the coordination by mercury as a fifth ligand. This weak interaction of CO to the adsorbed Ni(1) species for Ni(II)-TMC was confirmed by the experimental finding under CO that almost no positive potential shift was observed for the adsorption peak compared with the large positive shift for the Ni(II)/Ni(I)-TMC redox couple in solution. The spiky adsorption prewave for Ni(II)-TMC enabled us to confirm its negligible potential shift by CO in spite of the overlap with the shifted main wave, but confirmation was difficult for Ni(II)-cyclam with a broad prewave. Curves b and c in Fig. 10 show the theoretical current-potential curves for the cases where the concentrations of catalytically generated CO are assumed to be one-hundredth and one-tenth of that under CO at atmospheric pressure, respectively. In these calculations, the potential dependence of k,,, was neglected for brevity. The curve in Fig. 4 was used as the adsorption isotherm for calculation of the decrease in the catalytic current by the diminished concentration of Ni(II)-cyclam in the vicinity of the electrode surface. The qualitative resemblance between the experimental I-E curve and the calculated one indicates the validity of the interpretation of the decreased catalytic activity described above.
ACKNOWLEDGEMENTS
We would like to thank Dr. U. Akiba for helpful discussions and Mr. Y. Nakamura for assistance in part of the experimental work. This work was partially supported by a Grant-in-Aid for Scientific Research on Priority Area 63604534 and 63612506 from the Ministry of Education, Science and Culture.
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