Solution Redox Couples for Electrochemical Energy

Solution Redox Couples for Electrochemical Energy Storage I. Iron (lll)-Iron (11) Complexes with O-Phenanthroline and Related Ligands Yih-Wen D. Chen, K. S. V. Santhanam, z and Allen J. Bard* Department oJ Chemistry, The University of Texas at Austin, Austin, Texas 78713 ABSTRACT I r o n ( I I I ) - I r o n (II) complexes w i t h o - p h e n a n t h r o l i n e a n d r e l a t e d ligands have been e x a m i n e d b y electrochemical techniques in aqueous H2SO4 m e d i a w i t h respect to t h e i r s u i t a b i l i t y as r e d o x couples for electrochemical e n e r g y storage. The iron (II) complexes undergo a r a p i d 1 electron o x i d a t i o n at g r a p h i t e and p l a t i n u m electrodes to yield iron (III) complexes; these comp l e x e s showed v a r y i n g stabilities d e p e n d i n g on the n a t u r e of the substituents on the complexes. T h e iron (II) complexes e x a m i n e d in this s t u d y w e r e f o r m e d w i t h (i) monodentate, (ii) bidentate, or (iii) t r i d e n t a t e ligands. The r e d o x couples have a higher E o value w h i c h has been a positive consideration in the storage. A l t h o u g h the aquo iron ( I I ) - iron (III) couple has an E o less t h a n the complexes, it c e r t a i n l y has shown "greater p r o m i s e in terms of storage stability. The kinetics of iron (II) complexation~has been followed b y cyclic voltammetry. A r e d o x flow cell (or b a t t e r y ) (1-8) is o n e in which the chemical species which p a r t i c i p a t e in the electrode reactions are soluble. The cell is c h a r g e d w i t h the i n p u t of electrical e n e r g y to d r i v e the o v e r all cell reaction in a t h e r m o d y n a m i c a l l y u p h i l l d i r e c tion a n d the oxidized species p r o d u c e d at an i n e r t electrode in one h a l f - c e l l and the r e d u c e d f o r m f r o m the other are stored in e x t e r n a l vessels. E l e c t r i c i t y is p r o d u c e d in these cells w h e n the stored r e a c t a n t s flow back into the cell and r e a c t at the electrodes. Thus these cells are of i n t e r e s t as s e c o n d a r y or r e c h a r g e a b l e batteries. The e n e r g y densities of these systems (i.e., e n e r g y stored p e r unit w e i g h t of b a t t e r y ) suffer in c o m p a r i s o n to m o r e conventional seco n d a r y b a t t e r i e s which utilize solid reactants, b e cause of the weight of the solvent a n d electrolyte. However, t h e y offer the possibility of much b e t t e r cycle life, since the r e p e a t e d charge and discharge cycles do not involve phase changes a n d the a c c o m p a n y i n g changes of e l e c t r o d e morphology. These systems are of i n t e r e s t in s t a t i o n a r y applications such as electrical e n e r g y storage and in u t i l i t y load l e v e l ing. A n o t h e r r e l a t e d a r e a involving soluble r e d o x couples in e n e r g y devices is t h a t of p h o t o e l e c t r o chemical (PEC) (or liquid j u n c t i o n photovoltaic) cells w i t h semiconductor electrodes (9). These cells, w h i c h are based on the l i g h t - d r i v e n r e d o x processes of solution species at semiconductors, a r e of two types. In the PEC cell w i t h o u t e n e r g y storage, a single r e d o x couple is employed, and the electrode reaction at the counterelectrode is the r e v e r s e of the p h o t o - r e d o x process at the semiconductor. In PEC cells w i t h storage [types of p h o t o e l e c t r o s y n t h e t i c cells (9)], the r e a c t a n t s f o r m e d d u r i n g i r r a d i a t i o n are stored and e m p l o y e d to g e n e r a t e electricity, in the same or a s e p a r a t e cell, during d a r k periods. The r e d o x couples for these applications, r e p r e s e n t e d b y t h e reaction in [1] O ~ ne ~=~R ~ Electrochemical Society Active Member.

[1]

1Current address: Tata Institute of Fundamental Research. Bombay 400 005, India. Key words: battery, voltammetry, solubility, chelates.

m u s t satisfy a n u m b e r of r e q u i r e m e n t s : (i) both forms, O and R, m u s t be h i g h l y soluble to minimize the storage volume and mass and to allow high mass t r a n s f e r rates and c u r r e n t densities d u r i n g charging and discharging; (ii) the f o r m a l potential, E o', of one couple m u s t be h i g h l y positive, and E o' of the other h i g h l y negative to m a x i m i z e the cell voltage and e n e r g y density; (iii) the heterogeneous r e a c t i o n rate for the charging and discharging reactions at the i n e r t electrodes should be r a p i d (i.e., the s t a n d a r d r a t e constant, k o, for [1] should be large) so t h a t the electrode reactions occur at t h e i r mass t r a n s f e r controlled rates; (iv) b o t h O a n d R should be stable d u r i n g generation and storage, and this s t a b i l i t y p e r tains to reaction w i t h solvent, electrolyte, atmosphere, and electrode materials, and, for m e t a l complexes, s t a b i l i t y w i t h respect to l i g a n d loss; (v) the m a t e r i a l s should be safe, inexpensive, a n d a b u n d a n t ; (vi) the couple should not be corrosive and r e a c t w i t h cell materials, or the storage vessel, indeed, in PEC cells the r e d o x couple is often called u p o n to stabilize and protect the semiconductor electrode from photocorrosion (10, 11); and (vii) for PEC cells, the r e d o x species should not absorb light in the w a v e l e n g t h region of semiconductor absorption. A n u m b e r of r e d o x couples have been p r o p o s e d for such systems. T h e s e include F e ( I I I ) / F e ( I I ) (HC1) (5, 6); C r ( I I I ) / C r ( I I ) (HC1) (4-6); T i ( I V ) / T i ( I I I ) (7); Br2, B r - (12-14); and, for P E C cells, S 2-, Sx2 a n d Se 2-, Se22- (15, 16). A difficulty in a storage cell is the possible i n t e r m i x i n g of the components of the two half-cells, because of i m p e r f e c t s e p a r a t o r s or m e m b r a n e s , w h i c h leads not only to loss of cap a c i t y and efficiency b u t m o r e seriously to crossc o n t a m i n a t i o n of the r e d o x solutions. A p p r o a c h e s to minimizing this i n t e r m i x i n g p r o b l e m include the use of a single e l e m e n t system in t h r e e o x i d a t i o n states [e.g., C r ( V I ) , C r ( I I I ) / / C r ( I I I ) , C r ( I I ) (4)] or t h e use of two o x i d a t i o n states of the same e l e m e n t w i t h the r e d o x potentials shifted b y c o m p l e x a t i o n w i t h different ligands. The w o r k r e p o r t e d here, as w e l l as o t h e r c u r r e n t investigations (1, 2), is concerned w i t h the a p p l i c a t i o n of m e t a l ion coordination compounds as r e d o x

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Vol. 128, No. 7

1461

SOLUTION REDOX C O U P L E S

couples in flow cells. By suitable choice of ligand the formal potential of the couple can be shifted in the desired direction. Moreover such complexes may show improved characteristics with respect to stability and color in comparison to the uncomplexed species. The chemical principles related to the formation and properties of metal complexes are well developed and many new and potentially useful ligands have been reported. We describe here investigations of the F e ( I I I ) / F e ( I I ) couple with ligands related to ophenanthroline and bipyridine. Iron species were chosen for this initial investigation because iron is an abundant and inexpensive element with many highly soluble compounds. Such couples m a y be useful in the positive half-cell of redox flow systems or with

n-type semiconductors in P E C cells.

and an excess of ligand in 0.5M H2SO4. The complex formed would precipitate at this initial ion concentration. The precipitate was then just dissolved in the solution by adding 0.5M H2SO4 solution gradually. The concentration and the solubility were then estimated by the degree of the dilution of the solution (Table I). These measurements were done at room temperature. The stabilities of ferrous and ferric complexes were monitored with a Cary 14 spectrophotometer. The ferric complexes were prepared either by electrochemical oxidation of the ferrous complexes or by Ce (IV) oxidation in H2SO4 medium.

Results To shift the potential of the half-reaction

Experimental Reagents.--1,10-(or o - ) p h e n a n t h r o l i n e (phen), 2,9dimethylphenanthroline (2,9-dmp), 4,7-dimethylphenanthroline (4,7-drop), and 4-methylphenanthroline (mp) were o b t a i n e d ' f r o m Alfa Chemicals. 2,2'Bipyridine (bpy) was obtained from Aldrich Chemical Company. Tripyridine triazine (tpt), 4-cyanopyridine (cp), and 2-pyridine c a r b o x a l d e h y d e - 2 - p y ridyl hydrazine (p-cph)- were obtained from J. T. Baker Company. F e S O 4 . 7 H 2 0 was reagent grade (Matheson, Coleman and Bell). Surfactants I G E P A L Co430 (MW -- 396) and 530 (MW -- 484) were obtained from G A F Corporation (Atlanta, Georgia) and " T e x a s , l " (MW -- 404) was made by Dr. Y. B. Youssef of The University of Texas. The former two surfactants carry ethylene oxide groups and the latter is the sodium salt of 8 ( p - p h e n y l sulfanato) hexadecane. All solutions were prepared with doubly distilled water and the solutions were degassed with p r e p u r i fled gas that was passed through a chromous sulfate solution and then distilled water. Apparatus.--A Model 173 potentiostat in combination with a Model 179 digital coulometer (Princeton Applied Research Corporation, Princeton, New Jersey) was employed for all electrochemical experiments. The current-voltage curves were recorded on a Houston Instruments Model 2000 X - Y recorder. The current time curves were recorded on a Model 564 storage oscilloscope (Tektronix) during potential-step chronoamperometry and on a National Panasonic strip-chart recorder during coulometric experiments. Procedure.--The usual supporting electrolyte was aqueous H2SO4 prepared by suitable dilution of concentrated H2SO4. All solutions were degassed with nitrogen before the electrochemical experiments. The complexes were usually prepared directly in the electrochemical cell by mixing known concentrations of ferrous sulfate and the ligand. A mole ratio of l i g a n d / F e ( I I ) of greater than 5 was used. Controlled potential electrolyses were conducted with a large area graphite sheet electrode (area, 6.5 cm~) (Ultra Carbon, Sherman, Texas) with continuous nitrogen gas bubbling. Some of the electrochemical experiments were carried out in the dark. All of the potentials were measured with respect to an aqueous saturated calomel electrode (SCE). An H-cell with a porous sintered-glass disk separating the two compartments was used in coulometric investigations. For cyclic voltammetric investigations, a single compartment cell with a solution capacity of 5 ml was employed, with either platinum wire (A ---- 0.12 cm2), platinum disk (A -- 0.14 cm2), or graphite rod (taken from a C-cell battery, area A ---- 0.14 cm s) working electrodes. The platinum electrodes were pretreated by fast pulsing between +1.0 to --1.0V in H2SO4. The solubilities of the complexes were estimated by dissolving the specified concentration of FeSO4 9 7H20

F e 3+ + e :

[2]

F e ~+

toward values positive of the E o (+0.77V vs. NHE), ligands which form more stable complexes with Fe (II) are required. Since Fe 2+ (d 6) is a good ~-donor cation, ligands with low-lying vacant n* orbitals complex strongly with it (17). Fe ~§ has poorer ~donor properties because of its higher charge. Thus ligands such as bpy and phen (Fig. 1) are known to shift the potential of the redox couple in a positive direction. F u r t h e r manipulation of the potential is possible by substitution on the rings of these ligands. A number of highly stable complexes of these ligands are known (17). The structure and abbreviations for those ligands used in this study are shown in Fig. 1 and Table II. SolubiIities.--Table I is a list of the solubilities of the complexes in aqueous 0.5M H2SO4. Uncomplexed F e ( I I ) and F e ( I I I ) - s u l f a t e salts are quite soluble and yield solutions with metal ion concentrations ~ I M . The solubility of the complexes vary with the nature of the ligand. The ligands themselves are soluble in acidic media to ,.,2M (e.g., phen, 2,9-dmp). The complexes in 1M concentrations form solutions which are viscous and deeply colored. The uncomplexed F e ( I I ) and F e ( I I I ) ions in H2SO4 are quite transparent in the visible region.

E o' values.--The formal potentials of the various F e ( I I ) / F e ( I I I ) complexes were evaluated from the cyclic voltammetric peak potential values at low scan rates (Epa ----E ~ + 0.028/n) (see Table II). The complexes with cp, tpt, 2,9-drop, and p-cph show more positive peak potentials, but the reversible potential values are difficult to estimate because of the instability of the ferric complex. The Eo'-values for the other ligands are all more positive than E ~ for the aquo Fe ( 3 + / 2 + ) couple in H2SO4. In the phencomplexes substitution of methyl groups produces less positive peak potentials; while nitro- and chlorogroup substitution yields more positive potentials compared to complexes with the unsubstituted ligand. Similar observations have been made b y Fan and Table I. Estimated solubilities of ferrous-ferric couples and ligands*

Substance

Solubility (g/100 ml )

Concentratton (M)

FeSO~. 7I-I~O Fe2 ( SO~ ) 8-9H~O F e ( p h e n ) ~SO~ Fe ( b p y ) ~SO4 Fe (2,9-drop) SO4 F e ( tp ) ~SO~ Phenanthroline Bipyridyl Terpyridine 4-Cyanopyridine

27.8 44.0 52.6 43.8 37.2 22.4 40.0 31.6 44.0 10.5

1.0 0.8 0.8 0.7 1.0 0.3 2.2 2.0 1.8 1.0

* M e a s u r e d in 0.SM Ha$O4. See e x p e r i m e n t a l section f o r m e t h o d .

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J. Electrochem. Soc.: E L E C T R O C H E M I C A L S C I E N C E A N D T E C H N O L O G Y

1462

N~

C ~~-'~N

5 3 ?

6

~ 1

8. lO

9

Fig. i. Structure of ligands used in the study F a u l k n e r (18) electrode.

at a p h t h a l o c y a n i n e

covered m e t a l

Cyclic Voltammetric Measarements Electron transfer kinetics.--Cyclic v o l t a m m e t r y (cv) was e m p l o y e d to s u r v e y a system w i t h r e s p e c t to heterogeneous electron t r a n s f e r kinetics a n d s t a b i l i t y of the F e ( I I I ) - L species. Since the heterogeneous electron t r a n s f e r r a t e constants, k o, a r e all m o d e r a t e l y l a r g e ( > 1 0 - ~ c m / s e c ) , at low scan rates, v, the electrochemical b e h a v i o r (in the absence of fast following reactions) i s characteristic of a N e r n s t i a n r e a c t i o n and show AEp-values ,-,60 mV (AE : Epa -- Epc). w h e n v is increased, shifts of Epa and an increase in AEp a r e o b s e r v e d signaling the onset of effects of the heterogeneous electron t r a n s f e r r a t e (in the absence of u n c o m p e n s a t e d resistance effects) (19-21). I n f o r m a t i o n about ko can t h e n be o b t a i n e d from the v a r i a tion of Epa or AEp w i t h v using methods d e v e l o p e d b y M a t s u d a and A y a b e (19) a n d Nicholson (20). T h e a p p r o a c h t a k e n here was to d e t e r m i n e the v a l u e of r from the Ep-v b e h a v i o r (assuming a ~ / z ) , w h e r e ko

=

.'

C3]

X/~aDF~+ +

July 1981

diffusion coefficients of F e ( I I ) and F e ( I I I ) w e r e d e t e r m i n e d f r o m the p e a k c u r r e n t s ipa and ipc. F o r the aquo-species in 0.5M H2SO4, D F e 2 + - - D F e 3 + - - 2 . 5 X 10-6 cme/sec. T y p i c a l D - v a l u e s for the c o m p l e x e d F e ( I I ) species d e t e r m i n e d b y cv a r e listed in Table II. W i t h these D - v a l u e s and t h e m e a s u r e d ~-values, ko was obtained. These are also listed in T a b l e II. F o r example, in a t y p i c a l s t u d y w i t h 10 m M F e ( I I ) - p h e n c o m p l e x in 0.5M H2SO4, the cv w a v e showed Epa = 0.85V and, on scan reversal, Epc = 0.79V at v = 50 mV/sec. W h e n v increased, hEp i n creased, and at v = 50 V/sec, hEp -- 0.09V; v -- 100 V/sec, hEp = 0.10V; v : 200 V/sec, hEp -- 0.12V; v = 500 V/sec, AEp -- 0.18V. These m e a s u r e m e n t s w e r e m a d e w i t h full iR compensation. F r o m these values, an a v e r a g e value of k ~ -7- 5.8 • 10 - 2 c m / s e c was obtained. W i t h the F e ( I I ) coordination compounds only values for 0.5M H2SO4 a r e listed in Table II, since in 3M H2SO4 p r o t o n a t i o n of the ligands ( w i t h PKa's ~ 4-6) competes v e r y s t r o n g l y w i t h c o m p l e x a tion w i t h F e ( I I ) a n d a p p r e c i a b l e c o m p l e x a t i o n does not occur. Values of ko for the F e ( 3 + / 2 + ) couple in the a b sence of a d d e d l i g a n d w e r e also o b t a i n e d in,0.5 a n d 3M H2SO4, y i e l d i n g values of 1.6 X 10 - z a n d 1.2 X 10-2 cm/sec, respectively. T y p i c a l cv curves are shown in Fig. 2. The e l e c t r o c h e m i c a l o x i d a t i o n of F e (II) in aqueous H2SO4 has been p r e v i o u s l y i n v e s t i g a t e d at p l a t i n u m and carbon p a s t e electrodes (22, 23). The results are in g e n e r a l a g r e e m e n t w i t h those f o u n d here. T h e r a t e of the e l e c t r o d e reaction a p p e a r s to d e p e n d on the concentration of the acid as w e l l as the n a t u r e of the electrode surface (e.g., the presence of oxide films). However, the c o o r d i n a t e d species g e n e r a l l y have l a r g e r ko-values t h a n t h e aquo-species u n d e r similar conditions (0.5M H2SO4).

Stability (cv).--Stability of the coordination comp o u n d and the p r o d u c t of the e l e c t r o d e r e a c t i o n can also be m o n i t o r e d b y cv. Consider the o x i d a t i o n of FeLx 2+. The observed cv b e h a v i o r depends on the ligand, L, and the t e n d e n c y of FeLx s+ f o r m e d d u r i n g the anodic sweep to decompose (Fig. 3). D u r i n g the time scale of the cv sweep the F e ( I I I ) complexes w i t h p h e n (or b p y ) and tp a r e stable w h i l e t h a t w i t h cp is not. I n the ] a t t e r case oxidation of the stable F e ( I I ) c o m p l e x leads to t h e F e ( I I I ) form w h i c h d e composes rapidly, p r o b a b l y b y loss of l i g a n d to form a different F e ( I I I ) species. The g e n e r a l r e a c t i o n sequence is thus /co FeL~ s + r FeLx 3 + + e [4] kf FeL~3+ --> F e L x - ~ 8+ + y L

w h e r e a = n F v / R T a n d DFe+ + is the diffusion coefficient of t h e F e (II) ion in t h e m e d i u m employed. The

[5]

Table II. Thermodynamic and kinetic constants for ferrous-ferric couples* Ligands

E o' ( V vs. S C E )

Aquo

0.45

Bipyridine o-Phenanthroline 4-Methyl-o-phenanthroline 4,7-Dimethyl-o-phenanthroline 2,9-Dimethyl-o-phenanthroUne 5-Chloro-o-phenanthroline 6-Nitro-o-phenanthroline Terpyridine 2-P y r i d i n e e a r b o x a l d e h y d e

0.82 0.82 0.73 0.69 0.82 0.97w 1.{)7w 0.82 0.67 1.24 0.82

Tripyridinetriazine 4-Cyanopyridine

k~ (cm/sec)

1.6 1.2 6.6 5.8 4.9 2.5

• • • • • •

10-~ 10-2t 10-~ 10-2 10-2 10-2

6.0 x 10-2

Ir (sec -I x 10~)

0,063 0,012 0,006 0,003 >0.6 0.14 0,91 0.18 >0.6 > 0.6 >0.6

Ir ( M -~ sec -I • 10-4)

3,0 3.0 -6.3

800

D (cm ~ sec-: • 10 8)

Abbreviation of ligand

2.5

H~O

1.11 1.22 1.07 1.00

bpy phen mp 4,7.drop 2,9-drop C1P Np tp p-cph tpt ep

1.80

* M e d i u m 0.SM H~O~; e l e c t r o d e m a t e r i a l g r a p h i t e . T = 23 ~ "4- 2~ , , R a t e c o n s t a n t value refers to d e c o m p o s i t i o n of f e r r i c c o m p l e x e s d e t e r m i n e d b y s p e c t r o p h o t o m e t r y . T h e v a l u e ( > 0 . 6 ) w a s estim a t e d b y cv, t R a t e constant for t h e c o m p l e x a t i o n of f e r r o u s ion d e t e r m i n e d b y cv m e t h o d . $ M e d i u m 3M H~SO4. | K. Ogura and K. Miyamoto, ELeetroche~ica Acta, Z2, 1357 (1977).


VoI. 128, No. ?

SOLUTION REDOX COUPLES

1463

300 0

200

I (J.~)

0 C :D IVl Z

100 O

-I

0

-100

200

,

0.8

I 0.6

\ J ' OA

' 0.2

0.6

1. 20 mv/sec 2. 50 mv/sec 3 100 mv/sec

|C

(ma) 0.4

02

/ / f

0

-0,2

-0.4

E(u

t 0.4

I 0,2

I

I

I

I

1

0.8

0.6

0.4

G2

0

- 02

Fig. 3. Cyclic voltommetric curves for Fe(ll) ion coordinated with: 1, monodentate (cp); 2, bidcntate (bpy); 3, tridentate (tp) at graphite electrode. Solution contained 10 mM aquo Fe(ll) ion and at least 50 mM ligand, v = 0.05 V/sec.

0

Fig. 2. (a) Cyclic voltammetric curve for the oxidation of 10 mM Fe2+ in aqueous 3M H2S04 working electrode: Graphite, reference electrode: aqueous saturated calomel electrode, v = 0.50 V/sec.

~ 0.6

I

1.0

E ( V vs SCE )

300

E ( V ~s S C E )

0.8

[

1.2

|a

(.a)

-0,6 0

SCE)

Fig. 2. (b) The current-voltage curve after controlled potential exhaustive electrolysis at +0.80V in 0.SM H2S04.

(for simplicity, the state of profanation of the ligand is not indicated i n the equations). I n s t a b i l i t y o n the cv time scale for v up to 50 V/sec was also found for 2.9-dmp a n d several other ligands, as indicated i n Table II.

The rate of formation of the F e ( I I ) - l i g a n d complexes was also of interest and this was d e t e r m i n e d by cv for complexes with bpy, phen, 4,7-dmp, and tp. I n all eases the F e ( I I l ) complex is t h e r m o d y n a m i cally u n s t a b l e i n 0.5M H2SO4, although it is kinetically i n e r t so that it decomposes to Fe S+ a n d free ligand quite slowly. Thus if one prepares a solution containing a m i x t u r e of free Fe ~+ a n d L essentially no complexation occurs. Upon reduction of the Fe 3+ to F e 2+ rapid reaction with L takes place so that on scan reversal oxidation of F e ( l l ) occurs both as free Fe 2+ a n d as FeLx 2+, i.e., the reaction sequence is Fe~+ + e ~ Fe~+

[6]

Fe 2+ + zL-~ F e L ~ +

[7]

/co FeL~ +s ~ FeLx s+ + e

[8]

Typical results for a solution containing 3 mlVI Fe 3+ and 300 m M phen in 0.5M HzSO4 at a Pt electrode are shown in Fig. 4. Note that at very slow sweep rates (v -- 20 mV/sec) the anodic peak for Fe 2+ oxidation following Fe 3+ reduction is barely detectable, while the Fe(phen)32+ oxidation wave is fully developed. At higher sweep rates ipa for Fe 2+ oxidation increases and the Fe(phen)s ~+ oxidation wave becomes relatively smaller. At very high sweep rates (v -----100 V/see) only the Fe 2+ oxidation is observed with the same general characteristics as the uncomplexed Fe 3+ couple in the absence of phen at this v. The rate constant, kf' for Fe ~+ disappearance can be estimated from the ratio of ipa/ipc for the Fe 3+/2§ couple. Typical data for the cv of these complexes is shown in Table III along w i t h the rate constants obtained using the t r e a t m e n t of a secondorder following (EC) reaction (24). Since the ligand concentrations are high, the EC reaction is essentially pseudo-first order a n d the t r e a t m e n t of this E C - r e a c tion scheme (25) also applies. The calculated values of kf' are based on the free ligand concentration, ILl, at the given H2SO4 c o n c e n t r a t i o n which is i n e q u i l i b r i u m with the various protonated forms. The PKa values employed i n calculation of "[L] were:


J. EZectrochem. Soc.: E L E C T R O C H E M I C A L

1464

SCIENCE

AND

80

1.10 my/see 2.20 my/see 3.50 my/see 4.100 mv/eec 5.200 my/see

/,~ /

/

60

~

I

J u l y 1981

TECHNOLOGY FeL ~+ + L ~ F e L ~ 2 +

[10]

FeLs2+ + L ~ F e L ~ +

[11]

and reaction [8], where k o and k o' are heterogeneous e l e c t r o n t r a n s f e r r a t e c o n s t a n t s . I f [9] is t a k e n as t h e rate-determining step and neglecting dissociation rate, t h e s i m u l a t e d c u r v e s a r e n e a r l y i d e n t i c a l to t h e e x p e r i m e n t a l o n e as s h o w n i n Fig. 5. T h e s m a l l d i s c r e p a n c y f o r t h e F e 8+ r e d u c t i o n w a v e m i g h t b e d u e to t h e o n s e t of b a c k g r o u n d r e d u c t i o n . T h e s e c o n d o x i d a t i o n w a v e ( t h a t of F e L 3 2 + ) a n d i t s c o r r e s p o n d i n g r e v e r s a l r e d u c t i o n w a v e m i g h t also b e p e r t u r b e d b y t h e o c c u r r e n c e i n h o m o g e n e o u s s o l u t i o n of t h e r e a c tion k"

40

20

0

Fe~+ + FeLaS+ -, FeS+ + FeL82+

-20

-40

I 1.0

,

,

O~

, O~

~6

E( V vsSCE

-60

,

~2

0

)

Fig. 4. Cyclic voltammetric curve for 5 mM Fe(lll) reduction in the presence of phenanthroline (300 mM) at a platinum electrode,

b p y , p K = 4.47; p h e n , p K = 4.98; tp, p K ! - - 2.64, pK2 -~ 4.33 (26, 27). T h e k f ' - v a l u e s o b t a i n e d b y t h i s p r o c e d u r e a r e l i s t e d i n T a b l e II. A digital simulation of current-voltage curves was c a r r i e d o u t f o r t h e r e d u c t i o n of F e ( I I I ) i n t h e p r e s e n c e of t h e l i g a n d s . I n t h i s s i m u l a t i o n t h e f o l l o w i n g m e c h a n i s m w a s a s s u m e d . R e a c t i o n [6] w a s f o l l o w e d by y e 2 + + L ~ F e L 2+ [g]

[12]

A l t h o u g h t h e F e 2+ c o n c e n t r a t i o n i n t h e v i c i n i t y o f t h e e l e c t r o d e a t t h e p o t e n t i a l s of t h e s e w a v e s is s m a l l , t h i s r e a c t i o n is k n o w n to b e r a p i d , k" -- 2.2 • 105 1~ - 1 sec - 1 f o r L - - b p y i n 0.5M H2SO4 ( 2 8 ) . T h e i n c l u s i o n of t h i s r e a c t i o n i n t h e s i m u l a t i o n d i d n o t a f f e c t t h e o b s e r v e d i-E c u r v e s f o r k " - v a l u e s u p to 104 M - 1 s e c -1. W i t h i n t h e n u m b e r of i t e r a t i o n s u s e d i n t h e s i m u l a t i o n e x p l i c i t v a l u e s of k" h i g h e r t h a n t h i s c o u l d n o t b e e m p l o y e d . H o w e v e r , e v e n if k" is a s s u m e d to b e a t t h e m a s s t r a n s f e r c o n t r o l l e d l i m i t , so t h a t F e ~+ a n d F e L s 8+ c a n n o t c o e x i s t i n a s i m u l a t i o n s p a c e e l e m e n t , t h e e f f e c t o f t h e i-E c u r v e s , s h o w n as d a s h e d l i n e s i n Fig. 5, is s m a l l . I t is i n t e r e s t i n g to c o m p a r e t h e s e r e s u l t s o f c o m p l e x a t i o n of F e 2+ b y b p y w i t h t h o s e o b t a i n e d b y s t o p p e d flow m e t h o d s (29), w h e r e a s e c o n d - o r d e r

500 1. 2. 3. 4. 5.

10 mv/sec 20 mv/sec 50 m v / s e c 100 mv/sec 200 mv/sec

400 I

(pa) 300

Table III. Electrochemicalstudiesof complexation rate of Fe2+ : complexation with bipyridine,o-phenanthroline and tripyridine* 200 Sweep rate (V/see)

0.05 0.05 0.05 0.10 0.10 0.10 0.20 0.20 0.20 0.50 0.50

(pa/ipe

v (see)

kant (M-1 sec-~)

Concentration os bipyridine: 0.3M 0.230 5.00 4.00 0.275 4.00 3.51 0~254 4.00 4.11 0.307 2.50 4.52 0.367 2.00 3.94 0.341 2.00 4.66 0.435 1:25 4A9 0.539 1.00 3.65 0.458 1.00 5-06 0,619 0.50 5.54 0.656 0.40 6.60

k~$ (M-~ see-t • 10-~)

100 0.27 0.24 0.28 0.31 0.27 0.31 0.30 0.25 0.34 0.38 0.42

0.05

Concentration of o-phenanthroline: 0.3M 0.256 12,50 1.30 0.474 5.00 0.94 0.528 2.50 1.52 0.747 1.25 1.32 0.780 0.50 2.46 0.857 0.25 5.78 Concentration of terpyridine: O.02M 0.617 Ii.00 3.80

82.8

~.~5

Concentration of terpyridine: O.04M 0.416 11,20 4.11

75.8

0.02 0.05 O.lO 0.20 0.50 1.00

0.29 0.21 0.34 0.34 0.55 1,28

"Medium 0.5M H2SO4. Working electrode platinum. (Area = O.14 em~). Adsorption of the organic on the electrode sometimes produced irreproducibiUty. Temperature of the solutions ~25~ Concentration of Fe ~+ = 5.0 mM. t Calculated employing the total concentration of the ligand and assuming formation of the mona-c0mplex is rate determining. t Calculated with the effective ligand concentration uang pK, values reported in text.

0

-100

1D

~8

I ~6

I OA

I ~2

200 0

E (VvsSCE) Fig. 5. Digital simulation for the ECE process of the ferricbipyridine system in 0.SM H2S04. ( ) Experimental curves;digital simulation curves ( - - - ) as~suming infinite rate of reaction [12] and ( " . ) assuming rate of reaction [ 1 2 ] ----- 0. CFeS+ ~ 20 mM, Cbpy ---- 300 mM, kapp ---- 2.5 M - I s e e - l ; DFe3+ "2.5 X 10 - 6 cm2/sec, DFe(bpy)33+ ~ 2.6 X 10 - 6 cm2/sec. Area of platinum electrode ~ 0.148 cm 2. E~ 3+,2+) ~ 0.41V, E~ ~ 0.80V, a ( Fe3+ ' 2+ ) ~ 0.5, ~(Fe(bpy)3 3+ ,2+ ) ~ 0.5, ks(Fe 3+ ,2+ ) ~ 1.61 X 10 - 3 cm/sec, and ks(Fe(bpy)33+,2+) ~ 5.64 X 10 - 2 cm/sec.


Vol. 128, No. 7

1465

S O L U T I O N REDOX C O U P L E S

rate constant of 2.8 • 10~ M - I sec -1 at 0.2~ was reported. This rate constant was assigned to the first step in the complexation formation 9 of Fe(bpy) ~+ (reaction [9]). The addition of the second and third bpy ligands is faster; for the third step the rate constant is 1.4 X 105 M -1 sec -1, while for phen the rate of addition of the third ligand is 1.5 X 10e M -1 sec -1 (29). The medium and temperature effects on the dissociation rates of the bpy and phen complexes have been investigated by Basolo et al. (30).

100

Controlled Potential Coulometric (cpc) Measurements The long term stability and current efficiency for the processes on cycling were investigated by controlled potential electrolysis techniques. The oxidation of Fe e+ in the uncomplexed form carried out at 0.2V past the anodic c v peak showed an napp (faraday/mol) of 0.98-1.00. The current decayed smoothly to the background value and a plot of log i vs. t yielded a straight line (31) (Fig. 6). Reversal electrolYSiS, reduction of the electrogenerated Fe 8+, carried out at ,~0.10V, consumed essentially the same number of coulombs as in the forward electrolysis. These experiments were carried o u t at concentrations up to 1M and electrolysis times of ,-,0.5 hr. Oxidation of Fe(II)-complexes with bpy, phen, 2,9-dmp, and tp were carried out at 1.0V vs. SCE at a graphite electrode. In all cases the oxidations consumed 0.95-1.00 faraday/mol (see Table IV) and the current-time curve decayed smoothly to background. For electrolysis times ,~1.0 hr, reversal electrolysis of the solution at 0.10V consumed about the same number of coulombs as the forward electrolysis. Repetitive electrolysis, cycling between the oxidized and reduced forms four times at 0.5 hr intervals produced similar results. To examine the lifetime of the complex Fe(III) species, cv experiments were undertaken on solution following cpc oxidation. For ordinary cpc electrolysis of 30 min to 1 hr duration, cv showed reduction waves of height and location consistent with the presence of FeLx8+. However, when these solutions were allowed

100

2.0

40

LOG la

I a 30

(-) 1.0

2O

10

I

10

I

20

I

~

I

40

10

20

I9

1.0

LOG I c

30

40

2.0

F|g. 6. Current-time curve during exhaustive electrolysis of 10 mM Fe (phen)3 2+ in aqueous H2SO4 at a graphite working electrode.

50

0 I

50

0

-50

i

,

I

I

,

1.0

O~

O~

0,4

~2

-100 0

E (V vs SCE ) Fig. 7. Stability of Fe (phen)8 3+ in 1N H2SO4. Upper curve:

cyclic voltammetric curve for 10 mM Fe (phen)39+ immediately after controlled potential electrolysis at + I . 0 V at graphite electrode. Lower curve: cyclic voltammetric curve of the same solution after standing 24 hr in air.

to stand for times of~ 10-20 hr (Fig. 7), c v revealed that i~ for FeLxs+ reduction decreased and a wave for uncomplexed Fe 3+ appeared. Thus a slow decomposition of the FeLx8+ species (L ---- bpy, phen, rap, 4,7-dmp) does occur. The rate of decomposition of the FeLz8+ was determined by spectrophotometry, as described below. Stability of Fe(lll) Species Since the long duration storage capabilities of these systems appear limited by the stability of FeLx8+, measurements of the rate of decomposition of the Fe(III) species and some attempts at stabilization were carried out. The stability of Fe(III) complexes were examined spectrophotometrically. The Fe(!II) complexes were formed by chemical oxidation by mixing millimolar concentrations of Fe (II) complexes and excess Ce(IV) sulfate in 0.5M H2SO4 and monitoring the disappearance of the Fe(III) complex absorption (32). The rate constants for the first-order decay of the FeLx8+ species, k~, are listed in Table V. Substitution of a methyl group on phen more than doubles the half-life with 4,7-dmp being the most stable of the complexes examined. The 2,9-drop complex is very unstable, because in this case the placement of the methyl groups produces steric hindrance to bonding between the Fe(III) and the phenanthroline nitrogens. The background medium also appears to play a role in the rate of decomposition of Fe(III) complexes. Thus Fe(phen)88+ is reported to be stable in high concentrations of H2SO4 (22, 23). We found, in bulk!electrolysis experiments with the complexes, that the stability was decreased in H2804 solutions at a pH of -~1, compared with 0.5M H~SO4. The stability of Fe(II) and Fe(III) complexes has been studied previously (33, 34). The general conclusions are that the instability of the complex can be attributed to nucleophilic attack by water, resulting in replacement of the ligands by water molecule. The exact nature of


July 1981

J. Electrochem. Sac.: ELECTROCHEMICAL SCIENCE AND TECHNOLOGY

1466

Table IV. Controlled potential electrolysis of ferrous ions*

Moles taken

Q~o (C)

0.061 • 10-s 0.081 • 10-~ 1.19 x 10-~ 3.08 • 10-~ 3.00 x 10"~ 1.50 x 10-~ 3.00 X 10 -~

6.39 6.64 114,5 285.3 288.6

Moles of Fe s+ (oxdn)

Qb ~ (C)

Moles of Fe 2+

nappt (fara-

recovered (redn)

days tool-~)

6.3 x 10-5 6.5 • 10-~ 1.13 X 10-3 2.75 x 10 -s 2.78 x lO-a 1.52 • 10-~ 2.93 • 10-s

1.08 1.10 0:99 0.98 0.99 1.10 0,98

4.68 • 10-'

0.95

0.95 x 10"~

1.08

Free Fe ~+ions 0.006 • I0-~ 0.068 • 10-a 1.18 x 10"-s 2.96 x 10~ 2.99 x 10~ 1,70 x 10-2 2,93 • I0 "~

1708.9 2826.8

4.67 • 10-' 4.67 x 104

42.77

43.57

1,07 x 10-e

6.08 6,28 109,1 265.0 269.2

1470;0 2830.0

4.43 x 10~

Fe ( p h e n ) ~ + 45.20

1.07 x 10-~

92.00

4.51 x 10-~

104.10

41.00

4.24 X 10-4

0.97

Fe (2,9-dmp) z+ 3.10 x 10~ 3.10 x 10~

33.75

34.05

3.49 x 10-~

3.52 x 104

2.00 • 10-' 4.00 • i0-~

19.55

2.02 x 10-' 4.60 x 10"~

33.00 ~

3.4 x 10-' --

1.10 1.10

19.80 --

2.0 • 10-, --

1.02 1.10

Fe ( t p ) ~ 44.45

* W o r k i n g electrode: graphite. F o r w a r d electrolysis w a s carried out at +l.OV and reverse electrolysis at O.OV. t Values for oxidation.

the intermediate complex is not clear. Some authors (35) favor the reaction sequence FeL33+ -~ H 2 0 ~- FeL2 ( L . H 2 0 ) 3+ F e L 2 ( L 9 H 2 0 ) 8+ ~ F e I 4 ( L

9 O H ) 2+ + H +

[13] [14]

In our studies we observed the formation of free F e 3+ i n 0.5M H2SO4 f r o m F e L x ~+ u p o n s t a n d i n g f o r long duration, e v e n in the p r e s e n c e of excess ligand. This demonstrated that total replacement of the ligand b y H 2 0 e v e n t u a l l y o c c u r s , i.e. FeL83+ ~ F e e+ ~ 3L

[15]

( w h e r e L = p h e n , b p y ) . W h e n e x c e s s l i g a n d is n o t p r e s e n t , a c a t h o d i c p e a k a t -]-0.10V a p p e a r s o n s t a n d ing, w h i c h m i g h t r e p r e s e n t d i m e r i c s p e c i e s s u c h as a b r i d g e c o m p l e x f o r m e d (18) b y t h e f o l l o w i n g r e action 2FeL~ 8+ ~ 3 H 2 0 ~

Fe2L40 (H20)24+ -~ 2H + " t - 2 L

[16] Table V. Cyclic voltammetric data on the bulk electrolysis products*

Sweep rate Epa tps Epr /pc (V/see) (V) (mA) (V) (mA)

0.02 0.05 0.05

--

--

Remarks

0.05

0.48

5.20

0.05

0.86 0.86

Fe (phen)ar Cone = 20.0 mM 0.14 0.80 0.14 0.22 0.80 0.22

0.02

--

--

0.5

........ l,~

......... a ......

/f

0.80 0.87

Fe (2,9-drop) s+ 0.11 0.12 0.03 0.17 0.98 0.08

Cone = 20.0 m M B e f o r e b u l k electrolysis B e f o r e buLk electrolysis

0.10 0.11 After oxidation at + 1.0V 0.08

0.17

0.80 0.11 0.12 0.03 After reversal electrolysis at 0.17

Id

0.22

0.02 0.05

0,87

ik

0.80 0.14 After bulk electrolysis at + 1.0V

0.86 0.15 0.80 0.15 Reversal electrolysis at +0.26V 0,86 0,23 0.80 0.33

0.05

1.O

10

0.02 0.05

0.02

......... ~f

C

Reversal electrolysis at 0.0V

0.40

0.80

0.02 0.05

20 ~ ' ~ '

5.60

0.48 8.80 0.40

0.05

|k (.a}

0.40 3.90 After electrolysis at + 0.80V 0.40

0.02

0.02

s a t u r a t e d s o l u t i o n s of Li2SO4. T h e d e c a y o f t h e F e ( I I I ) complex concentration was noticed even in the prese n c e o f t h e a b o v e salts. W h e n s a t u r a t e d s o l u t i o n s w e r e used, the solutions b e c a m e viscous a n d the 30

Unbound Fe~+ Cone -- IM 0.48 3.80 0.40 Before bulk electrolysis 0.48 5.20 0.40 Before bulk electrolysis

0.02

Since attack by water causes decomposition of the Fe(III) species, blocking of the central metal ion by a hydrophobia environment or decreasing the water activity in the bulk solution may improve stability. Indeed the improved stability of the 4-mp and 4,7dmp may partially represent such an effect, although the electron-donating properties of the methyl groups may also play a role. Attempts at providing a more hydrophobic environment were made by introducing various surface active agents and monitoring the concentration of the Fe(III) complex by co. Results for the Fe(bpy)s 3+ complex are shown in Fig. 8 in 0.5M H2SO4 solutions containing butanol, heptanol, and the surfactant IGEPAL. Considerable improvement was noticed with 1-butanol in the medium. The abrupt changes in cyclic voltammetric peak current in the presence of a few surfactants (e.g., curve 8) occurs when phase separation takes place. Electrolysis of Fe(II) complexes was also conducted i n t h e p r e s e n c e o f 1M A1C18 o r 1M M g S O 4 o r

0,08

0.08

0.0V

* Measurements were made at graphite electrode in the same medium used for controlled potential electrolysis.

o

4

8

66 (,s)

Fig. 8. The decay of Fe(lll)-bpy complex in 1N H2S04 in the presence of stabilizing agents. The F e ( l l l ) - b p y complex was obtained by controlled potential electrolysis of Fe(ll)-bpy complex. /k = cyclic voltammetric peak current after controlled potential electrolysis (cathodic peak at -~0.80V) id is the diffusional peak Icurrent. i k = 1, 3, 5, 7, 9, and 11 and ik/id ~ 2, 4, 6, 8, 10, and 12, ik (t = 0) = 36 ~ A for line 7. I , 2 = No stabilizer; 3, 4 - Texas-1 (excess); 5, 6 = 5 ml 1-butanol; 7, 8 = IGEPAL Co-430; 9, 10 - - IGEPAL C0-530; 11, 12 = 1-heptanol.


Vol. 128, No. 7

SOLUTION REDOX C O U P L E S

duration of electrolysis was very long; with saturated solution in 0.5M H2SO4, electrolysis mmols Fe(phen)83+ took about 4 hr. At the electrolysis, a cathodic peak at O.31V for free was observed in the cyclic voltammograms.

LifSO4 of 12.6 end of Fe(III)

Discussion Aquo Fe(II)/Fe(III) couple in 6N H2SO4 appears to .be well suited as a redox positive couple for an extended period of time and appears to meet the major requirements of stability, solubility, fast electron exchange, and high electrical conductivity in the acid medium. A measurement of electron transfer rate constant in the 2N HC1 medium was undertaken to compare it with HeSO4 medium. At platinum disk or graphite electrodes, an average of five measurements gave k o values of 1.9 • 10 -3 and 1.34 X 10-8 cm/sec, respectively. These values are slightly smaller than the value obtained in H2SO4 (see Table I). However, a disadvantage of using HC1 medium is the somewhat less positive background limiting potential value compared with H2SO4 (22). An advantage of H2SO4 medium over Cl for PEC cells is the lighter colors of the Fe(II) and Fe(III) species in this medium. Complexation of F e ( I I ) / F e ( I I I ) couple can provide higher Eo'-values [e.g., 1.07V for np compared to 0.53V (vs. SCE)] and faster electrode kinetics. The solubility of the complexes is approximately the same as uncomplexed Fe(II) or Fe(III). While the Fe(II) complexes are highly stable over days in the absence of O~, Fe(III) complexes decay in the medium employed by loss of ligand (see Fig. 7). This process results in loss of stored energy since the redox couple so produced is at a less positive potential. However, this loss is slightly compensated by the fast complexation of Fe(II) by ligand during a discharge cycle. Note that if the rate of complexation of Fe(II) were rapid enough, the Fe(III) reduction wave would shift to potentials near those for the oxidation of the complexed Fe(II) species.

Conclusions The results here describe how coordination of the central metal ion can affect the potential and the kinetics of the electrode reaction to produce systems which might be utilizable in redox flow batteries. The aquo iron ( 3 + / 2 - 5 ) s y s t e m in H2SO4 appears to be an alternative to the iron (3-t-/2+) system in HC1 media currently under investigation. Complexation with phenanthroline or bipyridyl-type ligands results in significant positive shifts in the potential of the redox couple. The solubility of these couples appears satisfactory, but instability of the ferric form results in partial loss of capacity upon long-term storage.

Acknowledgment This work was supported by a grant from Texas Instruments, under TI/DOE cooperative agreement DE-FC01-79ER10,000. We appreciate the comments of a reviewer pointing out the significance of reaction [12] in the interpretation of the cv data. Manuscript submitted Oct. 20, 1980; revised manuscript received Feb. 20, 1981. Any discussion of this paper will appear in a Discussion Section to be published in the June 1982 JOURNAL. All discussions for the June 1982 Discussion Section should be submitted by Feb. 1, 1982.

Publication costs of this vaper were assisted by The University of Texas at Austin. REFERENCES 1. J. Giner, L. Swette, and A. Cahill, "Screening of Redox Couples and Electrode Materials" CR 134705, for NASA-Lewis Research Center, Cleve-

1467

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