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The electrochemical reduction of manganese dioxide in acidic solutions
J. Electroanal. Chem., 110 (1980) 145--158
145
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
THE E L E C T R O C H E M I C A L
R E D U C T I O N O F M A N G A N E S E DIOXIDE IN
ACIDIC SOLUTIONS P A R T III. V O L T A M M E T R I C
PEAK 3
J.A. LEE, W.C. MASKELL and F.L. TYE
Berec Group Ltd., Group Technical Centre, St. Ann's Road, London, N15 3TJ (England) (Received 19th October 1979)
ABSTRACT Linear sweep voltammetry in 0.0125--1.25 M H2SO4 on thin MnO2 films electrodeposited onto a smooth glassy carbon (GC) substarte reveals two cathodic current peaks. The first peak, peak 1, is terminated by loss of electrical conductivity of the deposit which remains essentially as 3"MnO2. Peak 2 occurs - 1 0 0 0 mV negative to peak 1 on a GC substrate. In 3.75--7.5 M HzSOa a further peak, peak 3, previously unreported in acidic solutions is observed on GC between peaks 1 and 2. Peak 3 arises according to the following cyclic scheme : ½MnO 2 + 1Mn2+ + 2 H + ¢- Mn 3+ + H20
Mn 3+ + e ~- Mn 2+
The reaction is instigated by the presence of small amounts of Mn 2+ generated previously during peak 1. The charge-transfer reaction proceeds on the GC substrate and not the MnO2 deposit as the latter has a very low electrical eonductivity after peak 1. Some thermodynamic measurements on H2SO 4 solutions are appended to supplement the general literature.
INTRODUCTION In a previous paper [1] the primary process in the electrochemical reduction o f t h i n f i l m s o f ~'-MnO2 e l e c t r o d e p o s i t e d o n t o p l a n a r P t , p l a t i n i s e d P t , p l a n a r A u a n d glassy c a r b o n ( G C ) s u b s t r a t e s w a s s t u d i e d i n H2 SO4 s o l u t i o n s u s i n g l i n e a r s w e e p v o l t a m m e t r y . I n 0 . 0 1 2 5 - - 1 . 2 5 M H~SO4 t w o c a t h o d i c c u r r e n t p e a k s w e r e o b s e r v e d . I n i t i a l l y t h e c u r r e n t r o s e i n t o p e a k 1, t h e r e a c t i o n proceeding on the MnO2 surface probably according to the scheme: M n O 2 + H ÷ + e -* M n O O H
MnOOH
+ H +-~ ~I M n 2+ + ~MnO2 + H 2 0
(1)
(2)
resulting in the overall reaction MnO2+4H
++2e-~Mn
2÷ ~ 2 H 2 0
(3)
However, during the electroreduction the electricalconductivity of the deposit decreased [I ] causing the current to diminish rapidly. Chemical and X-ray diffraction analyses of the remaining firmly adhering deposit showed it to be still
146
essentially 7-MnO2 between peaks 1 and 2. At more negative potentials peak 2 was observed, the process again proceeding according to the overall reaction (3) and ceasing upon complete removal of the MnO2 film. The position of peak 2, in contrast to that of peak 1, was dependent upon the composition of the substrate [2]. The potential separation between peaks 1 and 2 was greater on a GC than a platinum substrate. On raising the acid concentration into the range 3.75--7.5 M H2SO4 a third cathodic peak, peak 3, was observed between peaks 1 and 2 in the linear sweep voltammogram of an MnO2 deposit on GC. This paper concerns this new peak 3. EXPERIMENTAL
The techniques, procedures and equipment employed were as described previously [1] with the following additions. A 7.5 M H2SO4 solution approximately 0.02 M in Mn 3÷ was prepared by oxidising Mn 2÷ with permanganate. MnO~+4Mn 2÷+8H÷-~ 5Mn 3++4H20
(4)
The stoichiometric amount of 0.025 M KMnO4 was added dropwise to a vigorously stirred solution of MnSO4 in H2SO4while cooling in an ice bath [3]. The final volume was adjusted by adding water in the same manner. The solution was analysed by reducing with excess ferrous ammonium sulphate (-3 X 10 -3 M) and back titrating the excess Fe2+ potentiometrically with KMnO4 (~2.5 × 10-4 M). Continuation of the titration on the same solution after the addition of sodium pyrophosphate [4] enabled determination of the total Mn in the initial solution. This method was used by Vetter and Jaeger [5] for the analysis of small quantities of MnO2. Solutions of 7.5 M H2SO4both with and without Mn2+addition (~0.02 M) were also prepared. Mn2+/Mn3+/7.5 M H2SO4 solutions were made by mixing appropriate quantities of the stock solutions. All solutions containing Mn3+were stored at 5°C to inhibit oxygen evolution [6]. Liquid junction potential (ljp) measurements were made using the cell shown in Fig. 1. A hydrogen electrode (in zM H2SO4) was constructed which was interchangeable with the right-hand component of Fig. 1. The H2 used was commercial grade (99.9%) from British Oxygen and was not purified further. It was, however, pre-saturated with water vapour by bubbling through an H2SO4 solution of appropriate concentration. Potential measurements using the cell in Fig. 1 were made with a digital voltmeter (Bradley, Type 173B) except with solutions of very low conductivity where an electrometer (Electronic Instruments, Vibron) was used in conjunction with a precise voltage source (Bradley, dc voltage calibrator, Type 127). Potentials quoted of electrodes in H2SO4 solutions (zM) are the emf values of the cell, II
I[
HgJHg2SO4[0.5 M H2SO41114M KCII,Iz M H2SO4[MnO2 [I
II
Measurements were made without the salt bridge to avoid contamination of
147
REFERENCE f l ELECTRODE HglHg2SO~I0.SMH2SO~
i
TAP3 ~ i 2 : ~ O4
TAPI
05MH2S04 ~
A or
0,5MH2SO 4
A
B
Hg Hg2SO~I ,;/~HorKCL~IzM I H2SO/~IHg2SO/~Hg 0.5MH2S041 I]zMH2SO/.,jI
Fig. 1. Cell w i t h liquid j u n c t i o n s .
working solutions and corrected using data presented in the Appendix. Where reductions were carried out in Mn2+/Mna÷-containing solutions (zM H2SO4 + Mn2+/Mn3+) the lip correction applied was that appropriate to zM H2SO4. Where acidic pyrophosphate solutions were employed the potential shown is that of the cell, II
Hg IHg2SO410.5 M H:SO411110.5 M H~SO4 + 0.1 M Na4P20~IMnO2 II
Three types of MnO2 designated A, B and C were potentiostatically deposited from MnSO4/H2SO4 solutions at 90°C [1 ]. The overpotential for deposition increased in the order A < B < C. The type A deposit shows dendritic-type growth from a relatively small number of nucleation centres, while the type C deposit has a fibrous nature, the fibres emanating perpendicularly from the substrate from a large number of nucleation centres. The type B deposit shows properties intermediate between those of types A and C. Micrographs of the deposits have been presented previously [1]. In all cases the deposition charge was 500 mC cm -2. The GC substrate was mounted as a rotating-disc electrode in a leak-proof PTFE sheath [7]. Electrochemical reductions were performed at a sweep speed of 100 mV min-' and 25 ° C. RESULTS AND DISCUSSION
Potentiodynamic sweeps on stationary type A, B and C electrodes in 0.5 M H2SO4 are shown in Fig. 2. A large separation, ~1000 mV, between peaks 1 and 2 is evident and the current in the region between the peaks is very small. A cathodic voltammogram on a stationary type C electrode in 3.75 M H~SO4
148
0.8 0 .
0.6
0./.,
POTENTIAL I V 0.0
0.2
-0.2
~
-0.z,
-0.6
I
I
\
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-05
-0.8
I//i
\~/
V -I. 0 -
PEAKS 1
z~
~,J -I. 5 --
-2.0 --
_2sl Fig. 2. V o l t a m m o g r a m s o f d i f f e r e n t film t y p e s in unstirred 0.5 M H 2 S O 4 at 0.1 V m i n -1 o n GC : d e p o s i t 5 0 0 m C c m -2. (-- - - - - ) T y p e A; ( ) t y p e B; ( . . . . . . ) t y p e C. E l e c t r o d e area 0 . 2 0 c m 2 . POTENTIALI V 0,8 0.6
1.0 0
I
I
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0.0
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.
-0.2-
(a) ~
(b)
-0A-
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~
0
I
FLAKINGFROHSUBSTRATE
! \
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,
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(dl :
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I li IF z l vl~
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Fig. 3. V o l t a m m o g r a m s o f 5 0 0 m C c m -2 t y p e C e l e c t r o d e s in 3 . 7 5 M H 2 S O 4 at 0.1 V rain -1 o n GC. E l e c t r o d e r o t a t i o n s p e e d : ( ) stationary; (...... ) 200 rpm; ( ...... ) 1000 r p m . E l e c t r o d e area 0 . 2 4 c m 2 .
149 is s h o w n in Fig. 3a. T h e n o r m a l p e a k 1 o c c u r s first f o l l o w e d b y t w o o t h e r peaks. R o t a t i n g t h e e l e c t r o d e at 2 0 0 r p m (Fig. 3b) causes a d e c r e a s e in t h e m i d d l e p e a k a n d an increase in t h e p e a k at m o s t negative p o t e n t i a l s : this e f f e c t is progressive w i t h r o t a t i o n s p e e d (Fig. 3d). T h e p e a k at m o s t n e g a t i v e p o t e n tials is i d e n t i f i e d w i t h p e a k 2. T h e i n f l u e n c e o f e l e c t r o d e r o t a t i o n o n t h e m i d d l e p e a k , p e a k 3, i n d i c a t e s t h e p a r t i c i p a t i o n o f o n e or m o r e s o l u t i o n species. I t is d e m o n s t r a t e d in t h e s u b s e q u e n t discussion t h a t t h e m e c h a n i s m o f p e a k 3 is r e p r e s e n t e d b y t h e f o l l o w i n g cyclic s c h e m e : 1 MnO2 + ½Mn 2+ + 2 H ÷ ~- Mn 3÷ + H 2 0
(5)
Mn 3+ + e ~ Mn 2÷
(6)
R e a c t i o n (5) m a y involve t h e f o r m a t i o n o f M n ( I I I ) o n t h e solid MnO2 s u r f a c e b y an i n t e r a c t i o n w i t h Mn 2+ a n d s u b s e q u e n t d i s s o l u t i o n o f Mn3÷: 1 M N O 2 + 1~Ml~ _2+ + H 2 0 ~ - M n O O H + H
÷
(2)
M n O O H + 3 H + ~ Mn s+ + 2 H~O
(7)
A s w e e p w i t h i n t e r m i t t e n t stirring is s h o w n in Fig. 3e. W h e n r o t a t i o n is c o m m e n c e d d u r i n g p e a k 3 t h e c u r r e n t d e c r e a s e s r a p i d l y , t a k i n g u p a value similar t o t h a t d u r i n g c o n t i n u o u s r o t a t i o n . W h e n stirring is t e r m i n a t e d t h e curr e n t inereases again. Values o f (Q1 + Q2 + Q3)/QD (Q1, Q2, Qs a n d QD are t h e charges a s s o c i a t e d w i t h p e a k s 1, 2, 3 a n d d e p o s i t i o n r e s p e c t i v e l y ) f o r t h e s w e e p s o f Fig. 3a a n d 3e w e r e 0 . 9 2 a n d 0 . 8 9 r e s p e e t i v e l y . In 7.5 M H~SO4 l o w e r values w e r e r e c o r d e d (Table 1) i n d i c a t i n g t h e loss i n t o s o l u t i o n f r o m a t y p e C e l e c t r o d e o f m a n g a n e s e species o f v a l e n c y > 2 . I t was n o t e d t h a t u n d e r c e r t a i n e o n d i t i o n s a red c o l o r a t i o n was d i s c e r n i b l e in t h e s o l u t i o n i m m e d i a t e l y a d j a c e n t t o t h e e l e c t r o d e a n d t h e s e o b s e r v a t i o n s are s u m m a r i s e d in T a b l e 1. A c c o r d i n g t o R e m y [ 8 ] , M n 3÷ in e o n e e n t r a t e d H2SO4 has a red e o l o u r while Mn 4+ has a d e e p b r o w n e o l o u r in 5 0 - - 8 0 % H2SO4. T h e a b s o r p t i o n s p e c t r u m o f Mn 3÷ in a q u e o u s p e r e h l o r a t e [9] i n d i c a t e s t h a t t h e solut i o n s h o u l d a p p e a r red or r e d - p u r p l e . Visual c o m p a r i s o n w i t h t h e Mn3+/H2SO4
TABLE 1 Appearance of the red coloration. Electrode types -- A, B and C. Electrodes -- stationary. Colorations -- medium (M), faint (F) or not visible (N) Electrolyte
On open circuit
[H2SO4]/M
[MnSO4]/M
A
3.75 3.75 7.5 7.5 7.5
0 0.1 0 0.01 0.1
. --N F
During discharge
B .
C .
-----
. --M --
A . N N ---
(Q1 + Q2 + Q3)/QD
B
C
A
B
C
-M ---
0.97 0.97 0.95 0.97 0.96
0.96 0.95 0.92 0.96 0.90
0.92
N N F M
.
0.77 0.73
150
solution prepared by oxidation o f Mn 2÷ with MnO~ confirmed t h a t the coloration observed adjacent to the MnO2 was due to Mn 3÷. Trendafilov and Laitinen [10] detected soluble Mn(III) in H2SO4 solutions, both during the anodic formation o f MnO2 from Mn 2÷ and subsequent reaction (peak 1) using a rotating ring-disc electrode. Here, Mn 3÷ m a y be generated by the reaction o f Mn 2+ with MnO2 [11] : MnO2 + Mn 2÷ + 4 H ÷ -~ 2 Mn 3÷ + 2 H20
(5)
According to the following scheme [9] Mn a÷ is unstable in aqueous solutions: 2 Mn 3+ ~ Mn 4+ + Mn 2÷
(8)
--H +
Mn 4÷ + 2 H20 - " MnO2(solid) + 4 H ÷
(9)
+H +
Increasing Mn 2÷ or H + concentrations or complex formation tend to stabilise
Mn 3÷. The stability o f Mn 3÷ in H2SO4 is greatest at an acid concentration of 6.3 M [3] (half-life ~30 months). Selim and Lingane [3] measured equilibrium quotients for the disproportionation of Mn 3÷ in the range 4--7.2 M H2SO4, suggesting t h a t in the concentration range used in the present work and with the short intervals between the formation and reduction o f Mn 3* during linear sweep voltammetry of MnO2, the trivalent solution species m a y be considered as perfectly stable. From Table l the strengths o f coloration in the solution adjacent to the MnO2 vary as follows: electrode t y p e C > B > A; H2SO4 concentration 7.5 M > 3.75 M; MnSO4 concentration 0.1 M > 0.01 > 0. The (Q1 + Q2 + Q3)/QD values also decrease as the strength o f coloration increases due to the increase in the a m o u n t o f Mn 3÷ diffusing away into the bulk solution. After peak 1 the MnO2 deposit is a very poor electrical conductor [1] and reduction of the solution species evident as peak 3 must therefore take place on the GC substrate. Vetter and Manecke [6] studied the Mn2*/Mn 3÷ system in 7.5 M H2SO4 and concluded t h a t the charge-transfer reaction on Pt is simply Mn a+ + e ~ Mn 2÷
(6)
Cathodic current--voltage curves for the reduction of Mna÷/Mn :÷ solutions at a rotated GC electrode are shown in Fig. 4. At high overpotential (>100 mV) the current is controUed by charge transfer and diffusion resulting in linear plots of log[i/(1 -- i/il)] vs. potential as shown in Fig. 5: i is the current density and h the limiting value at high overpotentials. Within experimental accuracy the voltammograms in the Tafel region were independent of Mn 2÷ concentration and could be fitted to the equation [6] : i = --k3c3(1 -- i/il) exp[--(1 -- a) F E / R T ]
(10)
where k3 is the rate constant for the cathodic reaction, c3 the Mn 3+ concentration, ~ the transfer coefficient and E the electrode potential, the other symbols having their usual significance. Values of il were 152 and 1520 pA cm -2 in the 10 -3 and 10 -2 M Mn 3÷ solutions respectively. A linear regression analysis of the data in Fig. 5 indicated a value for a o f 0.69. This compares with the value o f 0.72 obtained by Vetter and Manecke [6] using a Pt electrode. The c o n f o r m i t y
151 [ '
,
,
i
,
,
]
I
I
I
i
3+ o f I0-3 10-2
e ,x I0-~"
I--
10-3 i
t
/./ Q
10-5
I
o O.O01MMn2÷
0:7
01.6 0!5
Oft* 01.3 0!2
POTENTIALI V
0
MI'I 2+'+
• O.01M i0-I,
0.9 0.8
0
0.1
IM~° 00 n t 0.7
0.6
÷
2
y I 0M 10.0 0.5
l O.t,
POTENTIAL / V
Fig. 4. Cathodic current--voltage curves for a GC electrode rotated at 1000 r p m in 7.5 M • . 2 + • and Mn 3 + ions.
H2S04 c o n t a m m g M n
Fig. 5. Tafel plots of cathodic current corrected for diffusion vs. potential for a GC electrode rotated at 1000 r p m in 7.5 M H2SO 4 c o n t a i n i n g - M n2÷ and Mn 3÷ ions.
b e t w e e n t h e results o n Pt [6] and GC indicates t h a t t h e p r i m a r y electron-transfer process o n b o t h substrates is t h e same, n a m e l y , Mn 3÷ + e ~ Mn 2÷
(6)
In a s t u d y o f t h e Mn3+/Mn 4* s y s t e m V e t t e r a n d M a n e c k e [12] s h o w e d t h a t a l t h o u g h t h e overall r e a c t i o n is Mn 4÷ + e ~ Mn 3+
(11)
t h e r o u t e is Mn 4+ + Mn 2÷ ~- 2 Mn 3÷
(8)
followed by Mn 3÷ + e ~ Mn 2÷
(6)
It was also a p p a r e n t [12] t h a t r e a c t i o n (8) is v e r y fast. In t h e p r e s e n t w o r k it is i n f e r r e d f r o m t h e s o l u t i o n c o l o r a t i o n t h a t t h e species f o r m e d in s o l u t i o n is Mn 3÷. Clearly, Mn 4÷ is also p r e s e n t d u e t o r e a c t i o n (8). H o w e v e r , t h e f o r m a l e q u i l i b r i u m q u o t i e n t Ks = [Mn3*]2/[Mn 4÷] [Mn 2÷] takes values [3] b e t w e e n 10 3 and 10 4 in 4--7.2 M H2SO4, indicating low Mn 4÷ concentrations. In t h e cyclic s c h e m e p r o p o s e d f o r p e a k 3 (reactions 5 and 6) if t h e r e were n o loss o f Mn 3÷ t o t h e s o l u t i o n , t w i c e as m u c h Mn 2* w o u l d be f o r m e d in r e a c t i o n (6) as originally i n t e r a c t e d with t h e MnO2 via r e a c t i o n (5). T h u s , t h e r e is a build-up o f M n 2+ and in t u r n M n 3÷ in t h e s o l u t i o n i m m e d i a t e l y a d j a c e n t t o t h e MnO2. T h e r e a c t i o n can t h e r e f o r e b e instigated b y t h e small a m o u n t o f Mn 2+
152
present in the solution generated during peak 1. The predominant solution species, Mn 2÷ and Mn ~÷, will tend to diffuse away from the electrode during peak 3. Suppose that a fraction g of the Mn 3÷ and a fraction h of the Mn 2÷ formed diffuse away. Then reaction (6) becomes (1 - - g ) Mn 3÷ + (1 - - g ) e ~- (1 - - g ) Mn 2+
(6a)
For a steady state (constant current) (1 - - g ) ( 1 - - h ) - 1
(12)
because ~Mn 2÷ is required b y reaction (5). If (1 - - g ) ( 1 -- h) >~ the current will increase until eqn. (12) is satisfied and vice versa. Peak 3 in Fig. 3a shows a maximum due to progressive depletion o f MnO:. If substantial amounts of Mn(III) in the solid plase were present after peak 1, then peak 3 in Fig. 3a could be explained b y the dissolution of this Mn(III) -~ Mn a÷ and reaction (6) only. Then, u p o n rotating the electrode, the Mn a÷ would tend to diffuse into the bulk solution and the peak 3 current would drop, as is observed. However, rotation would not also cause an increase in the height of peak 2 (compare Fig. 3a and 3d). Peak 2 represents the reduction to Mn 2÷ o f any MnO2 deposit remaining on the substrate below ~0.0 V (Fig. 2): the lack of rotation dependence (Fig. 11, ref. 1, Pt substrate) indicates that solution phase Mn 4÷ or Mn 3÷ are not primarily involved in the reaction during peak 2. It follows that Mn 3÷ is formed b y reaction (5) during peak 3 and that the deposit on the electrode must be essentially MnO2 after peak 1 as was demonstrated b y chemical analysis [ 1 ]. The latter part of a sweep on a t y p e B electrode in 3.75 M H2SO4 is shown in Fig. 6a. With no added Mn 2÷, peak 3 shows as a shallow plateau and peak 2 is prominent: the peak 3 current is smaller for t y p e B than for t y p e C electrodes. However, addition of 0.1 M MnSO4 increases the peak 3 current (Fig. 6a) and results in a small peak 2: this is compatible with the reaction mechanism proposed above (eqns. 5 and 6). Results for t y p e A electrodes in 3.75 M H2SO4 are shown in Fig. 6b. The effects noted when comparing t y p e B with t y p e C electrodes are even more pronounced with t y p e A. In 3.75 M H2SO4 peak 3 is absent, while in the presence of 0.1 M MnSO4 peak 3 appears as a shallow plateau. The differences in behaviour of t y p e A, B and C electrodes might be attributable to differing quantities o f deposit remaining after peak 1 (Fig. 13, ref. 1). However, this difference is small for t y p e A and B electrodes while the peak 3 behaviour is substantially different (cf. Fig. 6a and 6b). It is more likely to be a surface-area effect, the MnO2 area increasing in the order A < B < C. Scanning electron micrographs of the as-deposited films tend to support this idea (Figs. 4, 6 and 7 o f ref. 1), which also correlates with the strength of solution colorations noted earlier (A < B < C). The voltammogram o f a t y p e C electrode in 7.5 M H2SO4 is shown in Fig. 7a. Comparison with Fig. 3a indicates that increasing the acid concentration moves peak 3 to more positive values and leads to the disappearance of peak 2 (the MnO2 deposit completely removed before ~0.0 V). Rotation of the electrode (Fig. 7b) has a similar effect to that noted previously in 3.75 M H2SO4 and peak 2 is observed.
153 0.3
0.~ 0
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0.2
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-0.2
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//I
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0
0.4 I
0.3 I
0.2
01
0.0
I
L
i
i
I
(b)
-0.2
I
/ /
........<.\
-0.1 -
I
PEAKS 2
0.6
0.7
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0
0.5
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I
0.3
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/ ,.---. - ~ ~
-0.1-
(u
"\\
0.6
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0.0 I
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.OPEAKS2
""~ ._.j/ / "~PEAKS
0
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\\ -0.2
0.2
0.4 I
3
0.3 l
0.2 i
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-0 .I
-0.2 I
-0.3 i
-0.1-~
r
INFLECTIONS-PEAKS 2
' (d)
PEAKS 3
-0.2 4
POTENTIALI V
Fig. 6. V o l t a m m o g r a m s o f 5 0 0 m C c m -2 t y p e B a n d C e l e c t r o d e s at 0.1 V m i n -1 o n glassy c a r b o n : (a) t y p e B e l e c t r o d e in 3.75 M H2SO4; (b) t y p e A e l e c t r o d e in 3.75 M H2SO4; (c) t y p e B e l e c t r o d e in 7.5 M H2SO4; (d) t y p e A e l e c t r o d e in 7.5 M H2SO4. A d d i t i o n s t o t h e electrolyte: ( ) none; (..... ) 0.01 M M n S O 4 ; ( . . . . . . ) 0.1 M MnSO4. E l e c t r o d e area 0.24 cm 2.
154 1.0 0
0.8 I
0.6 I
POTENTIALIV 0.2 0.0 i NO PE;K 2
O.h I
-0.2 1
-O.Z+ I
-0.6
-0.2 E Z f~
PEAKS 3
I
I
1
I
~
I
.I
I
I
Fig. 7. V o l t a m m o g r a m s o f 500 mC c m -2 type C electrodes at 0.1 V rain -1 in 7.5 M H2SO 4. Electrode rotation speed: ( ) stationary; ( ...... ) 1 0 0 0 r p m . E l e c t r o d e a r e a 0 . 2 4 e m 2.
The behaviour o f t y p e B electrodes in 7.5 M H2SO4 is shown in Fig. 6c. As before, peak 3 is enhanced at the higher acid concentration (cf. Fig. 6a). The addition of MnSO4 increases the enhancement. Linear sweeps on t y p e A electrodes in 7.5 M H2SO4 are shown in Fig. 6d. Peak 3 is depressed compared with t y p e B or C electrodes, as was noted in 3.75 M H2SO4, and again MnSO4 enhances peak 3. In several cases shown in Figs. 6 and 7 peak 2 was absent. This was because in those instances all the deposit was removed from the GC substrate before ~0.0 V and the voltammograms in Fig. 2 indicate that peak 2 occurs at potentials negative to this value. The results m a y be summarised as follows: the reaction resulting in peak 3 is enhanced in the order: electrode t y p e A < B < C; H2SO4 concentration 3.75 M < 7.5 M; MnSO4 concentration 0 < 0.01 M < 0.1 M. The first effect m a y reflect surface-area differences while the other t w o are compatible with the proposed reaction scheme. The involvement of solution phase Mn 3÷ in the peak 3 reaction was confirmed b y complex ion stabilisation experiments. Here, Mn(III) forms a soluble complex with pyrophosphate in acidic solutions [4,9]. An electrolyte of 0.5 M H2SO4/0.1 M Na4P207 was used. The dissociation constants [4] o f H4P:O7 indicate that the predominant species in the above solution (pH ~ 0.7) is H3P20;. K o l t h o f f and Watters [13] found that a violet manganic pyrophosphate complex is formed in t h e pH range 3.8--0.1 and suggested that this is Mn(H2P2OT)~-. A linear potential sweep on a t y p e C electrode in a stationary pyrophosphate solution is shown in Fig. 8a. During the sweep a violet coloration in the solution adjacent to the electrode was noted, as indicated on the figure. The peak at more negative potentials was identified as peak 3, as electrode rotation
155
0.8
0.7
0.6
0.5
O.L,
o N -0,1
_o., <.,
VF
\/
AK 1
\
POTENTIAL / V 0.3 0.2
0.1
PEAK
0.0
-0-I
-0.2
-0.3
PEAK 2
/
~-o 1
-0.5 Fig. 8. V o l t a m m o g r a m s o f 5 0 0 m C e m -2 t y p e C e l e c t r o d e s in 0.5 M H2SO4/0.1 M N a 4 P 2 0 7 at 0.1 V m i n -1 • E l e c t r o d e r o t a t i o n s p e e d : ( ) stationary; (...... ) 200 rpm; Solution c o l o r a t i o n : (N) n o t visible; ( V F ) very f a i n t ; ( F ) faint. E l e c t r o d e area 0.24 e m 2.
produced a reduction in current (Fig. 8b). The suggested reaction mechanism is as follows: ~Mn02 + $M,I 1 _2+ + 3 H3P20~ ~ Mn(H2P2OT)~- + H20
(13)
Mn(H2P2OT)]- + 3 H+ + e ~ Mn2+ + 3 H3P20~
(14)
which is very similar to that proposed for the reaction at high acid strengths in the absence of pyrophosphate. According to Watters and Kolthoff [14], Mn2+ is predominantly present as Mn(H2P2OT)]- and eqns. (13) and (14) might be more correctly written as 1 2~MnO2+ ~Mn(H2P2OT)2
+ 2 H3P20~ ~ Mn(H2P2OT)~- + H20
Mn(H2P2OT)~- + H++ e ~- Mn(H2P2OT)~- + HaP20~
(15) (16)
CONCLUSION
Peak 3 results from a cyclic process involving the solution species Mn 2÷ and Mn 3+. The charge-transfer reaction occurs principally on the substrate. Peak 3 can be observed on GC because the peak 2 reaction generally takes place at more negative potentials on this substrate. ACKNOWLEDGEMENT
The authors wish to t h a n k the Directors of the Berec Group Ltd. for permission to publish this paper.
156
APPENDIX In this series of papers potential measurements of MnO2 were made using the cell A I I
Hg ]Hg2SO410.5 M H2SO4 z M H2SO4] MnO2] substrate
(17)
I
The true potential of the MnO2 vs. the standard hydrogen electrode (SHE) is given by (18)
F_, = E 1 7 + M -- Vz
E w is the emf of cell (w) and M is the potential of the left-hand electrode in cell (17) vs. SHE, not including the ljp Vz at A. The lip error Vz may be reduced by the introduction of a KC1 salt bridge: ]
I
Sg] Hg2SO4 [0.5 M I~I2SO414 M KCIIzM H2SO 4 [MnO2 ]substrate ]
(19)
I
Then E = E17 + M--S0.5 + Sz
(20) ]
Sz is the ljp of the junction z M H2SO414 M KC1. Equation (18) m a y be written: t
E, = E17 + ( M - - S o . 5 ) - - Vz + S0. s
(21)
With a knowledge of (M -- So. s) which is constant and (Vz -- So. s), E may be determined from E17 using eqn. (21). Potential measurements were made at 25°C using the following cells: l HglHg2SO410.5 M H2SO4tzM H2SO4 IHg2SO4 ]Hg (22) I
I
j
t
I
HglHg2SO4 i0.5 M H2SO4t 4 M KC1 Jz M H:SO41Hg2SO41Hg
(23)
]
Hg] Hg2SO 4 ]0.5 M H2SO41zM H2SOa ]H2 [Pt [ I
L
I
i
HglHg:SO410.5 M H2SO414 M KC1 IzM H:SO41H21Pt
(24) (25)
Cells (24) and (25) were constructed to cover the region z < 0.05 where the solubility of Hg2SO4 becomes a disturbing factor [15]. It can be seen that E22 -- E:3 = E2s -- E2,, = Vz --S0.s + S~
(26)
for any particular value of z. For low concentrations (z < 0.005), S~ is small and may be ignored. Then measurements of (E2s -- E24) provide values of (Vz - - S 0 . s ) according to eqn. (26). Measured values are presented in Table 2. In order to determine (M --S0.s) the cell I
PtIH2 ]zM H2SO414 M KClll Hg2Cl2 IHg I
I
(2 7)
157 TABLE 2 Voltages d e t e r m i n e d using cells (22)---(25)
z/M
0.0005 0.005 0.0125 0.05 0.05 0.2 0.5 1.25 3.75 7.5
E24/V
--0.7327 --0.7187 --0.7136 --0.7067 -------
E25/V
--0.8575 --0.7999 --0.7799 --0.7486 -------
E22/V
----0.047s 0.0197 0 --0.0234 --0.0714 --0.1293
E22/V (ref. 15) ----0.0483 0.0194 0 --0.0222 --0.0713 --0.1299
E23/V
(Vz -- S0.s + Sz)/V (eqn. 26)
----0.0054 0.001s 0 --0.0058 --0.0304 --0.0697
0.124s 0.0812 0.0663 0.0419 0.0424 0.0179 0 --0.0175 --0.041o --0.0596
was constructed. The right-hand electrode has a potential of 0.2459 V vs. SHE at 25°C including the ljp [16]. With z = 0.0005 the observed value of E27 was 0.426s. It may be shown t hat M -- S0.s = --E2s -- E27 + 0.2459
(28)
with E25 and E27 referring to the same z value. Applying eqn. (28) at z = 0.0005 leads t o M --S0.s = 0.677 V. This compares with the value 0.680 V q u o t e d elsewhere [12]. The small discrepancy can probably be traced t o different concentrations of KC1 used in t he salt bridge: this would influence the value o f So. s. Here, pH values m a y be calculated f r o m e m f values o f cell (25) as follows: E2s = ( R T / F ) In t~n+/UH2" /~1/2 -- Sz + So. s - - M
(29)
ax is the conventional activity of a species X disregarding hydration. A correction applied for atmospheric pressure [17] reduces all2 t o unity. T h e n pH = --(E2s + M --S0.s + Sz) F / 2 . 3 0 3 R T
(30)
Values o f E:4 for z > 0.05 may be d e t e r m i n e d by linear interpolation from the
TABLE 3 I n t e r p o l a t e d voltages f r o m refs. 15 a n d 18
z/M
E32/V
E24/V
E2 s / V
0.05 0.2 0.5 1.25 3.75 7.5
--0.7543 --0.719 s --0.6957 --0.6673 --0.6052 --0.521 s
--0.7060 --0.7001 --0.6957 --0.6901 --0.676 s --0.6514
--0.7484 --0.7180 --0.6957 --0.672s --0.635s --0.591 s
158 TABLE
4
pH vs. z
z/M
pH (eqn. 30) ignoring Sz
0.0005 0.005 0.0125 0.05 0.2 0.5 1.25 3.75 7.5 18.7
(conc.)
3.06 2.08 1.74 1.22 0.69 0.32 --0.068 --0.69 --1. 4 --
Literature values [19,20] -2.1 -1.2 -0.3 ----
--10
w o r k o f H a m e r [15] and H a r n e d a n d H a m e r [18] : E24 = E22 + E32
(31)
Hg IHg2 SO4 IzM H2 SO4 ]H21Pt
(32)
T h e e m f o f cell (25) can t h e n b e calculated f r o m E24 as s h o w n in Table 3 using eqn. (26) b y m a k i n g t h e a p p r o p r i a t e c o r r e c t i o n f o r (Vz --S0.s + S~) given in Table 2. Values f o r p H are p r e s e n t e d in Table 4 based o n E2s values in Tables 2 and 3 using eqn. (30) and ignoring Sz.
REFERENCES
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
J . A . L e e , W.C. M a s k e l l a n d F . L . T y e , J. E l e c t r o a n a l . C h e m . , 7 9 ( 1 9 7 7 ) 7 9 . P a r t II o f t h i s series. R . G . S e l i m a n d J . J . L i n g a n e , A n a l . C h i m . A c t a , 21 ( 1 9 5 9 ) 5 3 6 . J . J . L i n g a n e a n d R. K a r p l u s , I n d . E n g . C h e m . ( A n a l . E d . ) , 1 8 ( 1 9 4 6 ) 1 9 1 . K . J . V e t t e r a n d N. J a e g e r , E l e c t r o c h i m . A c t a , 11 ( 1 9 6 6 ) 4 0 1 . K . J . V e t t e r a n d G. M a n e c k e , Z. P h y s . C h e m . , 1 9 5 ( 1 9 5 0 ) 2 7 0 . W.C. Maskell, C h e m . I n s t r u m . , 8 ( 1 9 7 7 ) 4 3 . H . R e m y , T r e a t i s e o n I n o r g a n i c C h e m i s t r y , Vol. 2, Elsevier, A m s t e r d a m , 1 9 5 6 , p p . 2 2 4 , 2 4 6 . G. Davies, C o o r d . C h e m . R e v . , 4 ( 1 9 6 9 ) 1 9 9 . V. T r e n d a f i l o v a n d H . A . L a i t i n e n , M a n g a n e s e D i o x i d e S y m p o s i u m , V o l . 1~ E l e e t r o c h e m . S o c . , Clevel a n d , 1 9 7 5 , p. 7 9 . J.M. Cowley and A. Walkley, Nature, 161 (1948) 173. K . J . V e t t e r a n d G. M a n e c k e , Z. P h y s . C h e m . , 1 9 5 ( 1 9 5 0 ) 3 3 7 . I.M. K o l t h o f f a n d J . I . W a t t e r s , I n d . E n g . C h e m . ( A n a l . E d . ) , 1 5 ( 1 9 4 3 ) 8. J.I. W a t t e r s a n d I.M. K o l t h o f f , J. A m . C h e m . S o c . , 7 0 ( 1 9 4 8 ) 2 4 5 5 . W.J. Hamer, J. Am. Chem. Soc., 57 (1935) 662. R . G . B a t e s , D e t e r m i n a t i o n o f p H . T h e o r y a n d P r a c t i c e , Wiley, N e w Y o r k , 1 9 7 3 , p. 3 2 7 . G . J . Hills a n d D . J . G . I r e s i n D . J . G . I r e s a n d G . J . J a n z (Eds.), R e f e r e n c e E l e c t r o d e s , A c a d e m i c Press, N e w Y o r k , 1 9 6 1 , p. 9 6 . H . S . H a m e d a n d W . J . H a m e r , J. A m . C h e m . S o c . , 5 7 ( 1 9 3 5 ) 2 7 . Handbook of Chemistry and Physics, Chemical Rubber Publishing Co., Cleveland, OH, 41st ed., 1960, p. 1 7 2 1 . R . G . Bates, D e t e r m i n a t i o n o f p H . T h e o r y a n d P r a c t i c e , W i l e y , N e w Y o r k , 1 9 7 3 , p . 1 8 7 .
145
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
THE E L E C T R O C H E M I C A L
R E D U C T I O N O F M A N G A N E S E DIOXIDE IN
ACIDIC SOLUTIONS P A R T III. V O L T A M M E T R I C
PEAK 3
J.A. LEE, W.C. MASKELL and F.L. TYE
Berec Group Ltd., Group Technical Centre, St. Ann's Road, London, N15 3TJ (England) (Received 19th October 1979)
ABSTRACT Linear sweep voltammetry in 0.0125--1.25 M H2SO4 on thin MnO2 films electrodeposited onto a smooth glassy carbon (GC) substarte reveals two cathodic current peaks. The first peak, peak 1, is terminated by loss of electrical conductivity of the deposit which remains essentially as 3"MnO2. Peak 2 occurs - 1 0 0 0 mV negative to peak 1 on a GC substrate. In 3.75--7.5 M HzSOa a further peak, peak 3, previously unreported in acidic solutions is observed on GC between peaks 1 and 2. Peak 3 arises according to the following cyclic scheme : ½MnO 2 + 1Mn2+ + 2 H + ¢- Mn 3+ + H20
Mn 3+ + e ~- Mn 2+
The reaction is instigated by the presence of small amounts of Mn 2+ generated previously during peak 1. The charge-transfer reaction proceeds on the GC substrate and not the MnO2 deposit as the latter has a very low electrical eonductivity after peak 1. Some thermodynamic measurements on H2SO 4 solutions are appended to supplement the general literature.
INTRODUCTION In a previous paper [1] the primary process in the electrochemical reduction o f t h i n f i l m s o f ~'-MnO2 e l e c t r o d e p o s i t e d o n t o p l a n a r P t , p l a t i n i s e d P t , p l a n a r A u a n d glassy c a r b o n ( G C ) s u b s t r a t e s w a s s t u d i e d i n H2 SO4 s o l u t i o n s u s i n g l i n e a r s w e e p v o l t a m m e t r y . I n 0 . 0 1 2 5 - - 1 . 2 5 M H~SO4 t w o c a t h o d i c c u r r e n t p e a k s w e r e o b s e r v e d . I n i t i a l l y t h e c u r r e n t r o s e i n t o p e a k 1, t h e r e a c t i o n proceeding on the MnO2 surface probably according to the scheme: M n O 2 + H ÷ + e -* M n O O H
MnOOH
+ H +-~ ~I M n 2+ + ~MnO2 + H 2 0
(1)
(2)
resulting in the overall reaction MnO2+4H
++2e-~Mn
2÷ ~ 2 H 2 0
(3)
However, during the electroreduction the electricalconductivity of the deposit decreased [I ] causing the current to diminish rapidly. Chemical and X-ray diffraction analyses of the remaining firmly adhering deposit showed it to be still
146
essentially 7-MnO2 between peaks 1 and 2. At more negative potentials peak 2 was observed, the process again proceeding according to the overall reaction (3) and ceasing upon complete removal of the MnO2 film. The position of peak 2, in contrast to that of peak 1, was dependent upon the composition of the substrate [2]. The potential separation between peaks 1 and 2 was greater on a GC than a platinum substrate. On raising the acid concentration into the range 3.75--7.5 M H2SO4 a third cathodic peak, peak 3, was observed between peaks 1 and 2 in the linear sweep voltammogram of an MnO2 deposit on GC. This paper concerns this new peak 3. EXPERIMENTAL
The techniques, procedures and equipment employed were as described previously [1] with the following additions. A 7.5 M H2SO4 solution approximately 0.02 M in Mn 3÷ was prepared by oxidising Mn 2÷ with permanganate. MnO~+4Mn 2÷+8H÷-~ 5Mn 3++4H20
(4)
The stoichiometric amount of 0.025 M KMnO4 was added dropwise to a vigorously stirred solution of MnSO4 in H2SO4while cooling in an ice bath [3]. The final volume was adjusted by adding water in the same manner. The solution was analysed by reducing with excess ferrous ammonium sulphate (-3 X 10 -3 M) and back titrating the excess Fe2+ potentiometrically with KMnO4 (~2.5 × 10-4 M). Continuation of the titration on the same solution after the addition of sodium pyrophosphate [4] enabled determination of the total Mn in the initial solution. This method was used by Vetter and Jaeger [5] for the analysis of small quantities of MnO2. Solutions of 7.5 M H2SO4both with and without Mn2+addition (~0.02 M) were also prepared. Mn2+/Mn3+/7.5 M H2SO4 solutions were made by mixing appropriate quantities of the stock solutions. All solutions containing Mn3+were stored at 5°C to inhibit oxygen evolution [6]. Liquid junction potential (ljp) measurements were made using the cell shown in Fig. 1. A hydrogen electrode (in zM H2SO4) was constructed which was interchangeable with the right-hand component of Fig. 1. The H2 used was commercial grade (99.9%) from British Oxygen and was not purified further. It was, however, pre-saturated with water vapour by bubbling through an H2SO4 solution of appropriate concentration. Potential measurements using the cell in Fig. 1 were made with a digital voltmeter (Bradley, Type 173B) except with solutions of very low conductivity where an electrometer (Electronic Instruments, Vibron) was used in conjunction with a precise voltage source (Bradley, dc voltage calibrator, Type 127). Potentials quoted of electrodes in H2SO4 solutions (zM) are the emf values of the cell, II
I[
HgJHg2SO4[0.5 M H2SO41114M KCII,Iz M H2SO4[MnO2 [I
II
Measurements were made without the salt bridge to avoid contamination of
147
REFERENCE f l ELECTRODE HglHg2SO~I0.SMH2SO~
i
TAP3 ~ i 2 : ~ O4
TAPI
05MH2S04 ~
A or
0,5MH2SO 4
A
B
Hg Hg2SO~I ,;/~HorKCL~IzM I H2SO/~IHg2SO/~Hg 0.5MH2S041 I]zMH2SO/.,jI
Fig. 1. Cell w i t h liquid j u n c t i o n s .
working solutions and corrected using data presented in the Appendix. Where reductions were carried out in Mn2+/Mna÷-containing solutions (zM H2SO4 + Mn2+/Mn3+) the lip correction applied was that appropriate to zM H2SO4. Where acidic pyrophosphate solutions were employed the potential shown is that of the cell, II
Hg IHg2SO410.5 M H:SO411110.5 M H~SO4 + 0.1 M Na4P20~IMnO2 II
Three types of MnO2 designated A, B and C were potentiostatically deposited from MnSO4/H2SO4 solutions at 90°C [1 ]. The overpotential for deposition increased in the order A < B < C. The type A deposit shows dendritic-type growth from a relatively small number of nucleation centres, while the type C deposit has a fibrous nature, the fibres emanating perpendicularly from the substrate from a large number of nucleation centres. The type B deposit shows properties intermediate between those of types A and C. Micrographs of the deposits have been presented previously [1]. In all cases the deposition charge was 500 mC cm -2. The GC substrate was mounted as a rotating-disc electrode in a leak-proof PTFE sheath [7]. Electrochemical reductions were performed at a sweep speed of 100 mV min-' and 25 ° C. RESULTS AND DISCUSSION
Potentiodynamic sweeps on stationary type A, B and C electrodes in 0.5 M H2SO4 are shown in Fig. 2. A large separation, ~1000 mV, between peaks 1 and 2 is evident and the current in the region between the peaks is very small. A cathodic voltammogram on a stationary type C electrode in 3.75 M H~SO4
148
0.8 0 .
0.6
0./.,
POTENTIAL I V 0.0
0.2
-0.2
~
-0.z,
-0.6
I
I
\
/ /
\ PEAKS2
-05
-0.8
I//i
\~/
V -I. 0 -
PEAKS 1
z~
~,J -I. 5 --
-2.0 --
_2sl Fig. 2. V o l t a m m o g r a m s o f d i f f e r e n t film t y p e s in unstirred 0.5 M H 2 S O 4 at 0.1 V m i n -1 o n GC : d e p o s i t 5 0 0 m C c m -2. (-- - - - - ) T y p e A; ( ) t y p e B; ( . . . . . . ) t y p e C. E l e c t r o d e area 0 . 2 0 c m 2 . POTENTIALI V 0,8 0.6
1.0 0
I
I
0.2
0.A I
0.0
-0.2
I
.
-0.2-
(a) ~
(b)
-0A-
I
-0.6
I
~
0
I
FLAKINGFROHSUBSTRATE
! \
/ \
/
//
\ \ \
PEAKS3
\
\ \ \
!l / /
-0.8 I
I
/f
I~"~""FLAKINGSTARTED HERE
,
I "'':'~'" FLAKING
PEAKS2 i
(dl :
PEAKSI
F
I li IF z l vl~
I:
I. I. I1"
! I
I
I
I
I
I
Fig. 3. V o l t a m m o g r a m s o f 5 0 0 m C c m -2 t y p e C e l e c t r o d e s in 3 . 7 5 M H 2 S O 4 at 0.1 V rain -1 o n GC. E l e c t r o d e r o t a t i o n s p e e d : ( ) stationary; (...... ) 200 rpm; ( ...... ) 1000 r p m . E l e c t r o d e area 0 . 2 4 c m 2 .
149 is s h o w n in Fig. 3a. T h e n o r m a l p e a k 1 o c c u r s first f o l l o w e d b y t w o o t h e r peaks. R o t a t i n g t h e e l e c t r o d e at 2 0 0 r p m (Fig. 3b) causes a d e c r e a s e in t h e m i d d l e p e a k a n d an increase in t h e p e a k at m o s t negative p o t e n t i a l s : this e f f e c t is progressive w i t h r o t a t i o n s p e e d (Fig. 3d). T h e p e a k at m o s t n e g a t i v e p o t e n tials is i d e n t i f i e d w i t h p e a k 2. T h e i n f l u e n c e o f e l e c t r o d e r o t a t i o n o n t h e m i d d l e p e a k , p e a k 3, i n d i c a t e s t h e p a r t i c i p a t i o n o f o n e or m o r e s o l u t i o n species. I t is d e m o n s t r a t e d in t h e s u b s e q u e n t discussion t h a t t h e m e c h a n i s m o f p e a k 3 is r e p r e s e n t e d b y t h e f o l l o w i n g cyclic s c h e m e : 1 MnO2 + ½Mn 2+ + 2 H ÷ ~- Mn 3÷ + H 2 0
(5)
Mn 3+ + e ~ Mn 2÷
(6)
R e a c t i o n (5) m a y involve t h e f o r m a t i o n o f M n ( I I I ) o n t h e solid MnO2 s u r f a c e b y an i n t e r a c t i o n w i t h Mn 2+ a n d s u b s e q u e n t d i s s o l u t i o n o f Mn3÷: 1 M N O 2 + 1~Ml~ _2+ + H 2 0 ~ - M n O O H + H
÷
(2)
M n O O H + 3 H + ~ Mn s+ + 2 H~O
(7)
A s w e e p w i t h i n t e r m i t t e n t stirring is s h o w n in Fig. 3e. W h e n r o t a t i o n is c o m m e n c e d d u r i n g p e a k 3 t h e c u r r e n t d e c r e a s e s r a p i d l y , t a k i n g u p a value similar t o t h a t d u r i n g c o n t i n u o u s r o t a t i o n . W h e n stirring is t e r m i n a t e d t h e curr e n t inereases again. Values o f (Q1 + Q2 + Q3)/QD (Q1, Q2, Qs a n d QD are t h e charges a s s o c i a t e d w i t h p e a k s 1, 2, 3 a n d d e p o s i t i o n r e s p e c t i v e l y ) f o r t h e s w e e p s o f Fig. 3a a n d 3e w e r e 0 . 9 2 a n d 0 . 8 9 r e s p e e t i v e l y . In 7.5 M H~SO4 l o w e r values w e r e r e c o r d e d (Table 1) i n d i c a t i n g t h e loss i n t o s o l u t i o n f r o m a t y p e C e l e c t r o d e o f m a n g a n e s e species o f v a l e n c y > 2 . I t was n o t e d t h a t u n d e r c e r t a i n e o n d i t i o n s a red c o l o r a t i o n was d i s c e r n i b l e in t h e s o l u t i o n i m m e d i a t e l y a d j a c e n t t o t h e e l e c t r o d e a n d t h e s e o b s e r v a t i o n s are s u m m a r i s e d in T a b l e 1. A c c o r d i n g t o R e m y [ 8 ] , M n 3÷ in e o n e e n t r a t e d H2SO4 has a red e o l o u r while Mn 4+ has a d e e p b r o w n e o l o u r in 5 0 - - 8 0 % H2SO4. T h e a b s o r p t i o n s p e c t r u m o f Mn 3÷ in a q u e o u s p e r e h l o r a t e [9] i n d i c a t e s t h a t t h e solut i o n s h o u l d a p p e a r red or r e d - p u r p l e . Visual c o m p a r i s o n w i t h t h e Mn3+/H2SO4
TABLE 1 Appearance of the red coloration. Electrode types -- A, B and C. Electrodes -- stationary. Colorations -- medium (M), faint (F) or not visible (N) Electrolyte
On open circuit
[H2SO4]/M
[MnSO4]/M
A
3.75 3.75 7.5 7.5 7.5
0 0.1 0 0.01 0.1
. --N F
During discharge
B .
C .
-----
. --M --
A . N N ---
(Q1 + Q2 + Q3)/QD
B
C
A
B
C
-M ---
0.97 0.97 0.95 0.97 0.96
0.96 0.95 0.92 0.96 0.90
0.92
N N F M
.
0.77 0.73
150
solution prepared by oxidation o f Mn 2÷ with MnO~ confirmed t h a t the coloration observed adjacent to the MnO2 was due to Mn 3÷. Trendafilov and Laitinen [10] detected soluble Mn(III) in H2SO4 solutions, both during the anodic formation o f MnO2 from Mn 2÷ and subsequent reaction (peak 1) using a rotating ring-disc electrode. Here, Mn 3÷ m a y be generated by the reaction o f Mn 2+ with MnO2 [11] : MnO2 + Mn 2÷ + 4 H ÷ -~ 2 Mn 3÷ + 2 H20
(5)
According to the following scheme [9] Mn a÷ is unstable in aqueous solutions: 2 Mn 3+ ~ Mn 4+ + Mn 2÷
(8)
--H +
Mn 4÷ + 2 H20 - " MnO2(solid) + 4 H ÷
(9)
+H +
Increasing Mn 2÷ or H + concentrations or complex formation tend to stabilise
Mn 3÷. The stability o f Mn 3÷ in H2SO4 is greatest at an acid concentration of 6.3 M [3] (half-life ~30 months). Selim and Lingane [3] measured equilibrium quotients for the disproportionation of Mn 3÷ in the range 4--7.2 M H2SO4, suggesting t h a t in the concentration range used in the present work and with the short intervals between the formation and reduction o f Mn 3* during linear sweep voltammetry of MnO2, the trivalent solution species m a y be considered as perfectly stable. From Table l the strengths o f coloration in the solution adjacent to the MnO2 vary as follows: electrode t y p e C > B > A; H2SO4 concentration 7.5 M > 3.75 M; MnSO4 concentration 0.1 M > 0.01 > 0. The (Q1 + Q2 + Q3)/QD values also decrease as the strength o f coloration increases due to the increase in the a m o u n t o f Mn 3÷ diffusing away into the bulk solution. After peak 1 the MnO2 deposit is a very poor electrical conductor [1] and reduction of the solution species evident as peak 3 must therefore take place on the GC substrate. Vetter and Manecke [6] studied the Mn2*/Mn 3÷ system in 7.5 M H2SO4 and concluded t h a t the charge-transfer reaction on Pt is simply Mn a+ + e ~ Mn 2÷
(6)
Cathodic current--voltage curves for the reduction of Mna÷/Mn :÷ solutions at a rotated GC electrode are shown in Fig. 4. At high overpotential (>100 mV) the current is controUed by charge transfer and diffusion resulting in linear plots of log[i/(1 -- i/il)] vs. potential as shown in Fig. 5: i is the current density and h the limiting value at high overpotentials. Within experimental accuracy the voltammograms in the Tafel region were independent of Mn 2÷ concentration and could be fitted to the equation [6] : i = --k3c3(1 -- i/il) exp[--(1 -- a) F E / R T ]
(10)
where k3 is the rate constant for the cathodic reaction, c3 the Mn 3+ concentration, ~ the transfer coefficient and E the electrode potential, the other symbols having their usual significance. Values of il were 152 and 1520 pA cm -2 in the 10 -3 and 10 -2 M Mn 3÷ solutions respectively. A linear regression analysis of the data in Fig. 5 indicated a value for a o f 0.69. This compares with the value o f 0.72 obtained by Vetter and Manecke [6] using a Pt electrode. The c o n f o r m i t y
151 [ '
,
,
i
,
,
]
I
I
I
i
3+ o f I0-3 10-2
e ,x I0-~"
I--
10-3 i
t
/./ Q
10-5
I
o O.O01MMn2÷
0:7
01.6 0!5
Oft* 01.3 0!2
POTENTIALI V
0
MI'I 2+'+
• O.01M i0-I,
0.9 0.8
0
0.1
IM~° 00 n t 0.7
0.6
÷
2
y I 0M 10.0 0.5
l O.t,
POTENTIAL / V
Fig. 4. Cathodic current--voltage curves for a GC electrode rotated at 1000 r p m in 7.5 M • . 2 + • and Mn 3 + ions.
H2S04 c o n t a m m g M n
Fig. 5. Tafel plots of cathodic current corrected for diffusion vs. potential for a GC electrode rotated at 1000 r p m in 7.5 M H2SO 4 c o n t a i n i n g - M n2÷ and Mn 3÷ ions.
b e t w e e n t h e results o n Pt [6] and GC indicates t h a t t h e p r i m a r y electron-transfer process o n b o t h substrates is t h e same, n a m e l y , Mn 3÷ + e ~ Mn 2÷
(6)
In a s t u d y o f t h e Mn3+/Mn 4* s y s t e m V e t t e r a n d M a n e c k e [12] s h o w e d t h a t a l t h o u g h t h e overall r e a c t i o n is Mn 4÷ + e ~ Mn 3+
(11)
t h e r o u t e is Mn 4+ + Mn 2÷ ~- 2 Mn 3÷
(8)
followed by Mn 3÷ + e ~ Mn 2÷
(6)
It was also a p p a r e n t [12] t h a t r e a c t i o n (8) is v e r y fast. In t h e p r e s e n t w o r k it is i n f e r r e d f r o m t h e s o l u t i o n c o l o r a t i o n t h a t t h e species f o r m e d in s o l u t i o n is Mn 3÷. Clearly, Mn 4÷ is also p r e s e n t d u e t o r e a c t i o n (8). H o w e v e r , t h e f o r m a l e q u i l i b r i u m q u o t i e n t Ks = [Mn3*]2/[Mn 4÷] [Mn 2÷] takes values [3] b e t w e e n 10 3 and 10 4 in 4--7.2 M H2SO4, indicating low Mn 4÷ concentrations. In t h e cyclic s c h e m e p r o p o s e d f o r p e a k 3 (reactions 5 and 6) if t h e r e were n o loss o f Mn 3÷ t o t h e s o l u t i o n , t w i c e as m u c h Mn 2* w o u l d be f o r m e d in r e a c t i o n (6) as originally i n t e r a c t e d with t h e MnO2 via r e a c t i o n (5). T h u s , t h e r e is a build-up o f M n 2+ and in t u r n M n 3÷ in t h e s o l u t i o n i m m e d i a t e l y a d j a c e n t t o t h e MnO2. T h e r e a c t i o n can t h e r e f o r e b e instigated b y t h e small a m o u n t o f Mn 2+
152
present in the solution generated during peak 1. The predominant solution species, Mn 2÷ and Mn ~÷, will tend to diffuse away from the electrode during peak 3. Suppose that a fraction g of the Mn 3÷ and a fraction h of the Mn 2÷ formed diffuse away. Then reaction (6) becomes (1 - - g ) Mn 3÷ + (1 - - g ) e ~- (1 - - g ) Mn 2+
(6a)
For a steady state (constant current) (1 - - g ) ( 1 - - h ) - 1
(12)
because ~Mn 2÷ is required b y reaction (5). If (1 - - g ) ( 1 -- h) >~ the current will increase until eqn. (12) is satisfied and vice versa. Peak 3 in Fig. 3a shows a maximum due to progressive depletion o f MnO:. If substantial amounts of Mn(III) in the solid plase were present after peak 1, then peak 3 in Fig. 3a could be explained b y the dissolution of this Mn(III) -~ Mn a÷ and reaction (6) only. Then, u p o n rotating the electrode, the Mn a÷ would tend to diffuse into the bulk solution and the peak 3 current would drop, as is observed. However, rotation would not also cause an increase in the height of peak 2 (compare Fig. 3a and 3d). Peak 2 represents the reduction to Mn 2÷ o f any MnO2 deposit remaining on the substrate below ~0.0 V (Fig. 2): the lack of rotation dependence (Fig. 11, ref. 1, Pt substrate) indicates that solution phase Mn 4÷ or Mn 3÷ are not primarily involved in the reaction during peak 2. It follows that Mn 3÷ is formed b y reaction (5) during peak 3 and that the deposit on the electrode must be essentially MnO2 after peak 1 as was demonstrated b y chemical analysis [ 1 ]. The latter part of a sweep on a t y p e B electrode in 3.75 M H2SO4 is shown in Fig. 6a. With no added Mn 2÷, peak 3 shows as a shallow plateau and peak 2 is prominent: the peak 3 current is smaller for t y p e B than for t y p e C electrodes. However, addition of 0.1 M MnSO4 increases the peak 3 current (Fig. 6a) and results in a small peak 2: this is compatible with the reaction mechanism proposed above (eqns. 5 and 6). Results for t y p e A electrodes in 3.75 M H2SO4 are shown in Fig. 6b. The effects noted when comparing t y p e B with t y p e C electrodes are even more pronounced with t y p e A. In 3.75 M H2SO4 peak 3 is absent, while in the presence of 0.1 M MnSO4 peak 3 appears as a shallow plateau. The differences in behaviour of t y p e A, B and C electrodes might be attributable to differing quantities o f deposit remaining after peak 1 (Fig. 13, ref. 1). However, this difference is small for t y p e A and B electrodes while the peak 3 behaviour is substantially different (cf. Fig. 6a and 6b). It is more likely to be a surface-area effect, the MnO2 area increasing in the order A < B < C. Scanning electron micrographs of the as-deposited films tend to support this idea (Figs. 4, 6 and 7 o f ref. 1), which also correlates with the strength of solution colorations noted earlier (A < B < C). The voltammogram o f a t y p e C electrode in 7.5 M H2SO4 is shown in Fig. 7a. Comparison with Fig. 3a indicates that increasing the acid concentration moves peak 3 to more positive values and leads to the disappearance of peak 2 (the MnO2 deposit completely removed before ~0.0 V). Rotation of the electrode (Fig. 7b) has a similar effect to that noted previously in 3.75 M H2SO4 and peak 2 is observed.
153 0.3
0.~ 0
I
0.2
0.1
I
0.0
I
I
-0.I I
-0.2
-0.3
-0.4
-0.5
-0.6
-O-1
-0.2
-0.3
-O.L,
I
I
I
_~
I
//I
-0.1(Q)
PEA -0.2
0.6
0.5 i
0
0.4 I
0.3 I
0.2
01
0.0
I
L
i
i
I
(b)
-0.2
I
/ /
........<.\
-0.1 -
I
PEAKS 2
0.6
0.7
I
0
0.5
0.4
I
0.3
I
I
/ ,.---. - ~ ~
-0.1-
(u
"\\
0.6
0.7
0.5 i
0.1 ..L
/
0.0 I
-0.1 I
.OPEAKS2
""~ ._.j/ / "~PEAKS
0
I
/
\\ -0.2
0.2
0.4 I
3
0.3 l
0.2 i
0.1
0.0
I ___I___l
-0 .I
-0.2 I
-0.3 i
-0.1-~
r
INFLECTIONS-PEAKS 2
' (d)
PEAKS 3
-0.2 4
POTENTIALI V
Fig. 6. V o l t a m m o g r a m s o f 5 0 0 m C c m -2 t y p e B a n d C e l e c t r o d e s at 0.1 V m i n -1 o n glassy c a r b o n : (a) t y p e B e l e c t r o d e in 3.75 M H2SO4; (b) t y p e A e l e c t r o d e in 3.75 M H2SO4; (c) t y p e B e l e c t r o d e in 7.5 M H2SO4; (d) t y p e A e l e c t r o d e in 7.5 M H2SO4. A d d i t i o n s t o t h e electrolyte: ( ) none; (..... ) 0.01 M M n S O 4 ; ( . . . . . . ) 0.1 M MnSO4. E l e c t r o d e area 0.24 cm 2.
154 1.0 0
0.8 I
0.6 I
POTENTIALIV 0.2 0.0 i NO PE;K 2
O.h I
-0.2 1
-O.Z+ I
-0.6
-0.2 E Z f~
PEAKS 3
I
I
1
I
~
I
.I
I
I
Fig. 7. V o l t a m m o g r a m s o f 500 mC c m -2 type C electrodes at 0.1 V rain -1 in 7.5 M H2SO 4. Electrode rotation speed: ( ) stationary; ( ...... ) 1 0 0 0 r p m . E l e c t r o d e a r e a 0 . 2 4 e m 2.
The behaviour o f t y p e B electrodes in 7.5 M H2SO4 is shown in Fig. 6c. As before, peak 3 is enhanced at the higher acid concentration (cf. Fig. 6a). The addition of MnSO4 increases the enhancement. Linear sweeps on t y p e A electrodes in 7.5 M H2SO4 are shown in Fig. 6d. Peak 3 is depressed compared with t y p e B or C electrodes, as was noted in 3.75 M H2SO4, and again MnSO4 enhances peak 3. In several cases shown in Figs. 6 and 7 peak 2 was absent. This was because in those instances all the deposit was removed from the GC substrate before ~0.0 V and the voltammograms in Fig. 2 indicate that peak 2 occurs at potentials negative to this value. The results m a y be summarised as follows: the reaction resulting in peak 3 is enhanced in the order: electrode t y p e A < B < C; H2SO4 concentration 3.75 M < 7.5 M; MnSO4 concentration 0 < 0.01 M < 0.1 M. The first effect m a y reflect surface-area differences while the other t w o are compatible with the proposed reaction scheme. The involvement of solution phase Mn 3÷ in the peak 3 reaction was confirmed b y complex ion stabilisation experiments. Here, Mn(III) forms a soluble complex with pyrophosphate in acidic solutions [4,9]. An electrolyte of 0.5 M H2SO4/0.1 M Na4P207 was used. The dissociation constants [4] o f H4P:O7 indicate that the predominant species in the above solution (pH ~ 0.7) is H3P20;. K o l t h o f f and Watters [13] found that a violet manganic pyrophosphate complex is formed in t h e pH range 3.8--0.1 and suggested that this is Mn(H2P2OT)~-. A linear potential sweep on a t y p e C electrode in a stationary pyrophosphate solution is shown in Fig. 8a. During the sweep a violet coloration in the solution adjacent to the electrode was noted, as indicated on the figure. The peak at more negative potentials was identified as peak 3, as electrode rotation
155
0.8
0.7
0.6
0.5
O.L,
o N -0,1
_o., <.,
VF
\/
AK 1
\
POTENTIAL / V 0.3 0.2
0.1
PEAK
0.0
-0-I
-0.2
-0.3
PEAK 2
/
~-o 1
-0.5 Fig. 8. V o l t a m m o g r a m s o f 5 0 0 m C e m -2 t y p e C e l e c t r o d e s in 0.5 M H2SO4/0.1 M N a 4 P 2 0 7 at 0.1 V m i n -1 • E l e c t r o d e r o t a t i o n s p e e d : ( ) stationary; (...... ) 200 rpm; Solution c o l o r a t i o n : (N) n o t visible; ( V F ) very f a i n t ; ( F ) faint. E l e c t r o d e area 0.24 e m 2.
produced a reduction in current (Fig. 8b). The suggested reaction mechanism is as follows: ~Mn02 + $M,I 1 _2+ + 3 H3P20~ ~ Mn(H2P2OT)~- + H20
(13)
Mn(H2P2OT)]- + 3 H+ + e ~ Mn2+ + 3 H3P20~
(14)
which is very similar to that proposed for the reaction at high acid strengths in the absence of pyrophosphate. According to Watters and Kolthoff [14], Mn2+ is predominantly present as Mn(H2P2OT)]- and eqns. (13) and (14) might be more correctly written as 1 2~MnO2+ ~Mn(H2P2OT)2
+ 2 H3P20~ ~ Mn(H2P2OT)~- + H20
Mn(H2P2OT)~- + H++ e ~- Mn(H2P2OT)~- + HaP20~
(15) (16)
CONCLUSION
Peak 3 results from a cyclic process involving the solution species Mn 2÷ and Mn 3+. The charge-transfer reaction occurs principally on the substrate. Peak 3 can be observed on GC because the peak 2 reaction generally takes place at more negative potentials on this substrate. ACKNOWLEDGEMENT
The authors wish to t h a n k the Directors of the Berec Group Ltd. for permission to publish this paper.
156
APPENDIX In this series of papers potential measurements of MnO2 were made using the cell A I I
Hg ]Hg2SO410.5 M H2SO4 z M H2SO4] MnO2] substrate
(17)
I
The true potential of the MnO2 vs. the standard hydrogen electrode (SHE) is given by (18)
F_, = E 1 7 + M -- Vz
E w is the emf of cell (w) and M is the potential of the left-hand electrode in cell (17) vs. SHE, not including the ljp Vz at A. The lip error Vz may be reduced by the introduction of a KC1 salt bridge: ]
I
Sg] Hg2SO4 [0.5 M I~I2SO414 M KCIIzM H2SO 4 [MnO2 ]substrate ]
(19)
I
Then E = E17 + M--S0.5 + Sz
(20) ]
Sz is the ljp of the junction z M H2SO414 M KC1. Equation (18) m a y be written: t
E, = E17 + ( M - - S o . 5 ) - - Vz + S0. s
(21)
With a knowledge of (M -- So. s) which is constant and (Vz -- So. s), E may be determined from E17 using eqn. (21). Potential measurements were made at 25°C using the following cells: l HglHg2SO410.5 M H2SO4tzM H2SO4 IHg2SO4 ]Hg (22) I
I
j
t
I
HglHg2SO4 i0.5 M H2SO4t 4 M KC1 Jz M H:SO41Hg2SO41Hg
(23)
]
Hg] Hg2SO 4 ]0.5 M H2SO41zM H2SOa ]H2 [Pt [ I
L
I
i
HglHg:SO410.5 M H2SO414 M KC1 IzM H:SO41H21Pt
(24) (25)
Cells (24) and (25) were constructed to cover the region z < 0.05 where the solubility of Hg2SO4 becomes a disturbing factor [15]. It can be seen that E22 -- E:3 = E2s -- E2,, = Vz --S0.s + S~
(26)
for any particular value of z. For low concentrations (z < 0.005), S~ is small and may be ignored. Then measurements of (E2s -- E24) provide values of (Vz - - S 0 . s ) according to eqn. (26). Measured values are presented in Table 2. In order to determine (M --S0.s) the cell I
PtIH2 ]zM H2SO414 M KClll Hg2Cl2 IHg I
I
(2 7)
157 TABLE 2 Voltages d e t e r m i n e d using cells (22)---(25)
z/M
0.0005 0.005 0.0125 0.05 0.05 0.2 0.5 1.25 3.75 7.5
E24/V
--0.7327 --0.7187 --0.7136 --0.7067 -------
E25/V
--0.8575 --0.7999 --0.7799 --0.7486 -------
E22/V
----0.047s 0.0197 0 --0.0234 --0.0714 --0.1293
E22/V (ref. 15) ----0.0483 0.0194 0 --0.0222 --0.0713 --0.1299
E23/V
(Vz -- S0.s + Sz)/V (eqn. 26)
----0.0054 0.001s 0 --0.0058 --0.0304 --0.0697
0.124s 0.0812 0.0663 0.0419 0.0424 0.0179 0 --0.0175 --0.041o --0.0596
was constructed. The right-hand electrode has a potential of 0.2459 V vs. SHE at 25°C including the ljp [16]. With z = 0.0005 the observed value of E27 was 0.426s. It may be shown t hat M -- S0.s = --E2s -- E27 + 0.2459
(28)
with E25 and E27 referring to the same z value. Applying eqn. (28) at z = 0.0005 leads t o M --S0.s = 0.677 V. This compares with the value 0.680 V q u o t e d elsewhere [12]. The small discrepancy can probably be traced t o different concentrations of KC1 used in t he salt bridge: this would influence the value o f So. s. Here, pH values m a y be calculated f r o m e m f values o f cell (25) as follows: E2s = ( R T / F ) In t~n+/UH2" /~1/2 -- Sz + So. s - - M
(29)
ax is the conventional activity of a species X disregarding hydration. A correction applied for atmospheric pressure [17] reduces all2 t o unity. T h e n pH = --(E2s + M --S0.s + Sz) F / 2 . 3 0 3 R T
(30)
Values o f E:4 for z > 0.05 may be d e t e r m i n e d by linear interpolation from the
TABLE 3 I n t e r p o l a t e d voltages f r o m refs. 15 a n d 18
z/M
E32/V
E24/V
E2 s / V
0.05 0.2 0.5 1.25 3.75 7.5
--0.7543 --0.719 s --0.6957 --0.6673 --0.6052 --0.521 s
--0.7060 --0.7001 --0.6957 --0.6901 --0.676 s --0.6514
--0.7484 --0.7180 --0.6957 --0.672s --0.635s --0.591 s
158 TABLE
4
pH vs. z
z/M
pH (eqn. 30) ignoring Sz
0.0005 0.005 0.0125 0.05 0.2 0.5 1.25 3.75 7.5 18.7
(conc.)
3.06 2.08 1.74 1.22 0.69 0.32 --0.068 --0.69 --1. 4 --
Literature values [19,20] -2.1 -1.2 -0.3 ----
--10
w o r k o f H a m e r [15] and H a r n e d a n d H a m e r [18] : E24 = E22 + E32
(31)
Hg IHg2 SO4 IzM H2 SO4 ]H21Pt
(32)
T h e e m f o f cell (25) can t h e n b e calculated f r o m E24 as s h o w n in Table 3 using eqn. (26) b y m a k i n g t h e a p p r o p r i a t e c o r r e c t i o n f o r (Vz --S0.s + S~) given in Table 2. Values f o r p H are p r e s e n t e d in Table 4 based o n E2s values in Tables 2 and 3 using eqn. (30) and ignoring Sz.
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