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Oxidation of some aromatic amines on the passivated hanging mercury drop electrode
J. Electko@n@.
-Chek;
-0 El~evier Sequ@ia__S;A.,
65’(1975) 651-659 Lausanne - Printed~ in Thk Netherlands
.OXPDATiON OF SOME AROtiATIC AMINES MAhTGING MERCURY DROP ELECTRODE*
WIRTOR
KEMULA,
BARBARA
kstitute
of Physical
Chemistry,
(Received Zlst
July
BEHR Poiish
and JOANNA Academy
ON THE PASSIVATED
TARASZEWSKA
of Sciences, Wursaw
(Poland)
1975)
ABSTRACT It: has been shown that a hanging mercury drop electrode of the Kemula type is a very convenient tool for studying the redox processes of organic substances in the positive potektial range. Copditions influencing the oxidation process are discussed_ The method described-may prove to be an analytical tool for the determination of some aromatic amines.
3INTRODUCTION
Electrooxidations of organic substances in the far anodic potential range ha& recently received considerable attention. A lot of work in this field has been done on the Pt, Au, graphite and carbon paste electrodes. The mercury pool electrode covered with a fii of calomel was used successfully for organic oxidations by Kuwana and Adams [I] as well as by Podbi&nc& [2] _ -4 hanging tiercury drop electrode of the Kern&a type [ 33 is mu&more-advantageous, because.it can be easily reproduced in the same soiution.-- Uzkg this electrode pa&vated in chloride or bromide sol&ions we were able [ 43 toeoxidize.ferrocyanide and ferrous ions. F_ormation of slightly soluble mercurou& salts in the course of anodic dissolution of mercury in the presence of sm& c¢&i&s of Cl-, Br- and I- ions was used by us for the analytical. determination of these anions [ 5 3 _ The.possibility application of the passivated han&ng~mercury drop electrode fPH$J.DE) for studying redox processes in the positive potential range ‘-markedly depexids~on the kind of the-compound to be oxidized and @so on th.e properties of -the passive layer tid the rnamjef. of its formation, as well &s ,on: the tiompositi6.n of tiltition in ‘which p_assivation of :mercury and subse_ ’ qu_e%oxidat&‘of the brg&c mofec&es take piace. -_
of
652
Fig. 1. Currentvoltage curves recorded at the hanging mercury drop electrode 800 mY min-l in 0.1: M solutions of KCZ, KBr, KF, KI, KOH, and K,SO,.
at scan i-ate
Our previous St&dies [6] concerning the character and properties of the passive layer formed on the Hg electrode in solutions of various anions showed that the -kind af anion is of fundamental IGnportance for the behaviour of Hg electrode (see Fig. 1). From various experiments we concluded that the Hg electrode covered with a layer of Hg2Cls or HgzBr2 should be the most promising for study of redox processes in the positive potential range, In studies of the initial stages of electrooxidation and passivation of Hg in KC1 solutions f7] we have found that the mechanism of film formation as well as the thickness &.nd structure of the blocking film are dependent on the concentration of ;KCl. In dilute KC1 solutions (0.1-0.2 M) the blocking film is rather thick (on the average PO-20 molecules) and porous with an ohmic resistance of the order of 10 S2 cm2. In more concentrated solutions the film is compact and in 0.5 M &Cl it may be only about 3 molecules thick. The aim of this paper was to use the PHMDE in chloride solutions for the study of the redox processes of some aromatic amines and to find the best conditions of film formation fos this purpose. EXPERIMENTAL
A hanging mercury drop electrode with a s&ace area-of (0.02 2 0.002) cm2 was used- as previously described [3] _ Chronovoftamperometric curves were ‘-recorded using a. Radiometer polarograph PO& ~411.solutions were deaerated wit$ electrolytically generated hydrogen. The supporting electrolyte was 0.1 bM %iClI cf. &&.lyt&al reagent grade. In several cases purifi* methanol wris add-ed ir?order~to improve the soltibility: of the-studibd compounds; All ‘organic substances -empioy.& were additionally purifiedby re&y-stallisation or dist$l~ .-_
-~ _
5
653
lation. A calomel electrode (0.1 M KCI) was used as reference. In some cases the charge~consumed in oxidation and reduction processes was measured with a microcoulometer @3] of a special type (polaroquanter)3 The passive film was always formed in the presence of the oxidizable species. RESULTS
The HMDE was polarized anodically in solutions of various concentrations of CI- ions and pH at scan rates 100, 200, 400 and 800 mV min-I. The optimum conditions for the subsequent anodic oxidation of the organic substances were found to be in solution of 0.1 M HCl at scan rate 800 mV min-‘. Measurements for solutions containing only the supporting electrolyte are shown as curve 1 on the following diagrams_ Peak (a) corresponds to the formation of calomel. The origin of peak (a’) is not well understood but it seems likely to be connected with some further reactions of mercury with Cl- ions. The presence of methyl alcohol- in solution had no significant influence on the possibility of oxidation, but as we have found previously 193 other anions such as NOT and ClO, act destructively on the passive layer and may not be tolerated above a concentration exceeding l/10 that of the Cl- ions.
of
Oxidation-reduction
benzidine
Results for experiments with benzidine are shown in Fig. 2. Peak (b) corresponds to the oxidation of benzidine. In the concentration range 2 X lo-’ -
+oJ3
.-
‘a4
E/y
Fig. 2. Ckrentholtage
tions of benzidine;
a’ curves
for 0.1
(1) 0, (2) 5 X.10m4
171 HCl solution_ containing the following concentraM, (3) l.0s3 M, (4) 2 5 10” M, (5) 4 X 10m3 M_
Fii. 3. Cyclic current-voltage
curves for 0.1 M HC‘I containing
10H3 M benzidine.
102M t.he height of peak (b) is directly proportional to the concentration of benzidine in the bulk solution. If the direction of the voltage sweep is reversed after the peak of benzidine oxidation is attained a corresponding cathodic peak appears (peak (c), curve 1, Fig. 3). The charges consumed in the benzidine oxidation and the subsequent reduction of the product formed were measured coulometrically with the polaroquanter and the ratio of charges was found to be 3: 1 when the cathodic curve was recorded immediately after registration of the oxidation peak. A multiple cyclic polarization under these conditions only slightly changes the heights of pe_ak (b) and (c) (curve 2, Fig_ 3): However, polarization of the e&&rode to more positive potentials makes the prccess highly irreversible
--Fig. 4. q>qent---volwe
CURVES for Q.l~M He1 sdlution containing (A) o-tohidine in co% cenGatiotir (I) 0, (2) 2 X 1UI3 Lx, (3) 4 _XI.0m3 M, (4) 8 X 10d3 ,?I, (.5) lO+ M,-(B) ptoluidine in cqncentrations: (I) 0, (2) LOW3 M, (3) 2 X 10H3 M, (4) 4 X 10V3 n/Ii (5) 3.0-a+ _
655
(curve 3, E’ig. 3) leading to. the compfete disappearance cf the cathodic peak (curve 4, Fig, 3). Thephenomena describedare similartothoseobsemed previously by us 14-j in the course -of oxidation of ferroeyanide on the PEIMDE, Oxidation of other amines Oxidation curves for isomeric o- and p-toluidines, p-aminoacetophenone, N-methylaniline, p-phenylenediamine and diphenyfamine are shown in Figs. 4-8, respectively. Oxidation of toluidines, p-aminoacetophenone and N-methylaniline occurs at rather positive potentials near +I V. In this potential region the electronic conductivity of the passive layer is poor and the oxidation currents were small. Furthermore, no &n-rents corresponding to reduction of the oxidation prodttcts formed could he observed. In the case of p-aminoacetophenone we found that the oxidation peaks were higher from solutions containing 10 ~01% of CHaOH than from aqueous solutions. Attempts to effect oxidation of m-toluidine were unsuccessful. Results for F-phenyfenediamine and diphenylamine were similar to those of benzidine. Initial oxidation of p-phenylenediamine occurs at a potential of
-+-+ I -i_ f
-z--
_!_
---I-+-I--k I -y--
+
L
7
5. Current-voltage curves for 0.1 fM HC1 + 10 vd% MeOH containing p-aminoacetophenone in concentrations: (I) 0, (2) lo-’ M, (3) 2 X low3 M, (4) 4 X low3 N, (5) 6 X 10v3 M, (6) 10--2&L Pig_
Fig. 6. Ckren~oltage curves for 0.1 -&IHCL containing N-methylaniline in concentrations: Q) 0, (2).5-X IO--~ N. (3) IO--? M, (4) 2 x 10-S M, (-5) 4 X 1O-3 AI, (6) 1O-2 I&
Fig. 7. Current-voltage tions: (1) 0, (2) 10e3 Fig- 8. Current-voltage in concentrations: (1)
M,
curves for 0.1 M HCl containing p-phenylenediarnine (3) 2 X lo-' M, (G%) 4 X 10M3 A.2.(5) 6 X 10e3 M_
in concentra-
curves for 0.1 .&I HCI -P-50 ~01% CH,OH containing diphenyiamine 0, (2) 2 x low4 kf, (3) 4 X 10B4 &f, (4) 6 X 10e4 1M, (5) low3 BE
+0.5 V at which the second peak in pure chloride solution also appears- Additional peaks at higher positive potentials shown in Fig. 7 indicate that the inG tial product undergoes further reactions. Oxidation of diphenylamine was studied in a solution containing 50 ~019%C;H30K, because of its low solubility. Again in the cases of both p-phenylenediamine and diphenylamme no reduction of the formed oxidation products co&d be observed. DfSCUSSION
The results showed that the Kemula type hanging mercury drop electrode [3], passivated in a solution of Cl- ions, can be a very convenient tool for studying the electrooxidation of organic compounds in the anodic potential rangeIn the studies by Kuwana and Adams Cl] and by Podbielancik [SJ of the oxidation of organic compounds, they used a passivated, mercury pool efectrade_-Podbielancik’s attempts to use the hanging mercury drop electrode of &he Gerischer type [lOI were unsuccessful- The-use of the mercury pool is a rather tedious technique because in order to obtain good reproducibility each experiment requires the-renewal of the solution and of the mercury. :Moreover in-both papers -[I J and [Zj it was-recommended that the organic substance should be added after the formation of the passive layer.
657
Our present experiments indicate similarly to the case of the ferrocyanide oxidation 141 that the presence of the oxidizable substances in solution before formation-of a passivating film has no visible influence on~the possibility of oxidation. It follows from our experiments that the possibility of organic electrooxidation largely depends on the structure and properties of the passive layer as well as on the kind of substituents and their position iy a molecule.
The most suitable experimental &angement for our purposes was the Hh!IDE covered with a rather thick and porous passive layer formed in dilute HCl solutions. In 0.1 M HC1 the aromatic amines are mainly in the form of cations with the ratio [RNHs]‘J[RNH,] > 103. The significance of the interaction of the -ir-electrons of the aromatic ring with the Hg surface for the process of adsorption of aromatic compounds at potentials corresponding to the positive branch of the electrocapillary curve has been shown by Frumkin [11] and subsequently by Gerovich and Rybal’ chenko LIZ]. In addition Blomgren and Bockris {13] as well as Conway and Bmdas [14] have stressed the significance of r-electron interaction with Hg in the adsorption process of various aromatic a-mines from acidic and neutral solutions. They have found that in acidic media the amine cations should lie flat on the Hg surface at potentials corresponding to the positive branch of the electrocapillary curve. We believe that when the mercury electrode is covered with calomel the interactions of rr-electrons of the aromatic ring with the passive layer should also play an important role in oxidation process. This view is supported by the fact that we were able to oxidize only aromatic amines. Attempts at oxidation of some aliphatic compounds were unsuccessful. If the oxidation reactions were diffusion controlled the plots of peak current vs- concentration should be linear with a slope in agreement with the RandlesdevCik equation [15] _ We only get such a dependence in the case of benzidine, where we assumed n = 2 and D = 5 X 10v cm2 s-l. The reversible oxidation of benzidine to diiminequinone was proved by Mat&a et aI- 11161 at TABLE
1
Comparisonof E1,s and Ep values at Pt and PHMDE,
respectively,
measured vs. 0.1 NCE
atpH=l
Compound
+JV
Eknzidine o-Toluidine PToluidine N-niethylaniline pAminoacetophenone Diphetiyiarnine p_Phenylenediamine
+0.580 +0.99 - io.93 +x.075 -4-l-02 +0.93 - +0.83 +1.05 +0.69 +0.52
%21V
0.57 0.82 0.79
[16] [lS) [IS]
_-.@jf3.
-. y-_
-_-
:
_- _-:
i 2
4
6
Fig. 9. Calibration Fig.
_ ._.
_.-
_- =
-
IO. Calibration
8
f0
10%/M curves
curves
for N-methylaniline for benzidine
(I),
(I),
p-toluidine
diphenylamine
(2)
and o-toluidine
(3).
(2) and p-phenylenediamine
(3).
the Pt disc electrode in a wide pH range (6-9):There is good agreement between the oxidation potentials at the Pt and PHMBE which suggest that the mechanism of benxidine oxidation is the same at both electrodesA comp_arison of other results is not relevant because the reactions are Meversible and-E, values depend on experimental conditions. Anodic oxidation of a series bf p-substituted anilines in aqueous solutions was studied at the carbon electrode by 3acon and Adams [17]_ They found a common mecha-nis_minvolving elimination of the para groups arid h&ad-to&ii coupling. The corresponding 4 substituted 4aminodiphenylamines were formed as oxidation products_ A sititiar mechanism may be postulated in the case- of the I p-substituted aniline-oxidation at *he PHMDE. Our experiments have.shown that only the ortl20and para-isomers could be oxid@d. Difficulties with the oxidation of_.~~eta:isomers are dtie to the lack of formation of the quinone strticturein this lattercase.The tialibration curves for the -compounds -which we have studied
..
.:
__
:
-.. _~ 1
-
_- ;-
__
..
659
~-. . .
3 W.:_Kemulti&d 4 W: Ke-tiula, %
Z. tiLb!&, Anal; Chti. &&, 18 (l&,8) I&. Ktiblik and- d.~Taraszewska, J. EIectroanaJ. Chem., 6._(1963) 119. 5 *_ Ke@xuila,-2. J&blik-and J_ Taras&&vska, Proc:I~t. Symp_ on Microchemical Tech: ~n@ues, (19_Sl) 865: _ ----6 W. K&mula~and~J_ Taraszewske, Re;; Chim. Minerale, 5 (1968) 535. -7 B. Behr and J:~Taraszewska, J_ EJectroam& Chem., 19 (1968) 373. 8 -J-W: Strtijek and J. Zak, J_ Electroan& Chem., 46 (1973) 435. 9 -2. Kublik and J_ Terasze+&ka, BuII. Acad_ PoJon_ Sci_ Ser. Sci. Chim., 10 (1962) 515. 10 R. Gerischer, Z. Phys. Chem., 20 (1953) 302. 11 A-N. Frumkin; Ergebn. exakt. Naturwiss., 7 (1928) 256. 12 M.A. Gerovich and G.F. Rybal’chenko, Zh. Fiz. Khim., 32 (1958) 109. 13 E. Biomgren and J.O’M. Bockris, J. Phys. Chem., 63 (1959) 1475. 14 B.E:.Conway and R.G. Barradas, Electrochim. Acta, 5 (z961) 319, 349. 327; A SevEik, Collect. Czech. Chem. 15 J.E_B. Randles, Trans. Faraday Sot., 44 (1948) .Commun., 13 (1948) 349_ 16 M. Matrka, J. Pipaloti, Z. Sdgner and J. Marhold, ChemP&ml., 21 (1971) 14. 17 J_ Bacon and R.N. Adams, J. Amer. Chem. Sot., 90 (1968) 6596. 18 P-G. Desideri, L, Lepri and D. Heimler, J. Electroanal. Chem., 52 (1974) 93, 105.
-Chek;
-0 El~evier Sequ@ia__S;A.,
65’(1975) 651-659 Lausanne - Printed~ in Thk Netherlands
.OXPDATiON OF SOME AROtiATIC AMINES MAhTGING MERCURY DROP ELECTRODE*
WIRTOR
KEMULA,
BARBARA
kstitute
of Physical
Chemistry,
(Received Zlst
July
BEHR Poiish
and JOANNA Academy
ON THE PASSIVATED
TARASZEWSKA
of Sciences, Wursaw
(Poland)
1975)
ABSTRACT It: has been shown that a hanging mercury drop electrode of the Kemula type is a very convenient tool for studying the redox processes of organic substances in the positive potektial range. Copditions influencing the oxidation process are discussed_ The method described-may prove to be an analytical tool for the determination of some aromatic amines.
3INTRODUCTION
Electrooxidations of organic substances in the far anodic potential range ha& recently received considerable attention. A lot of work in this field has been done on the Pt, Au, graphite and carbon paste electrodes. The mercury pool electrode covered with a fii of calomel was used successfully for organic oxidations by Kuwana and Adams [I] as well as by Podbi&nc& [2] _ -4 hanging tiercury drop electrode of the Kern&a type [ 33 is mu&more-advantageous, because.it can be easily reproduced in the same soiution.-- Uzkg this electrode pa&vated in chloride or bromide sol&ions we were able [ 43 toeoxidize.ferrocyanide and ferrous ions. F_ormation of slightly soluble mercurou& salts in the course of anodic dissolution of mercury in the presence of sm& c¢&i&s of Cl-, Br- and I- ions was used by us for the analytical. determination of these anions [ 5 3 _ The.possibility application of the passivated han&ng~mercury drop electrode fPH$J.DE) for studying redox processes in the positive potential range ‘-markedly depexids~on the kind of the-compound to be oxidized and @so on th.e properties of -the passive layer tid the rnamjef. of its formation, as well &s ,on: the tiompositi6.n of tiltition in ‘which p_assivation of :mercury and subse_ ’ qu_e%oxidat&‘of the brg&c mofec&es take piace. -_
of
652
Fig. 1. Currentvoltage curves recorded at the hanging mercury drop electrode 800 mY min-l in 0.1: M solutions of KCZ, KBr, KF, KI, KOH, and K,SO,.
at scan i-ate
Our previous St&dies [6] concerning the character and properties of the passive layer formed on the Hg electrode in solutions of various anions showed that the -kind af anion is of fundamental IGnportance for the behaviour of Hg electrode (see Fig. 1). From various experiments we concluded that the Hg electrode covered with a layer of Hg2Cls or HgzBr2 should be the most promising for study of redox processes in the positive potential range, In studies of the initial stages of electrooxidation and passivation of Hg in KC1 solutions f7] we have found that the mechanism of film formation as well as the thickness &.nd structure of the blocking film are dependent on the concentration of ;KCl. In dilute KC1 solutions (0.1-0.2 M) the blocking film is rather thick (on the average PO-20 molecules) and porous with an ohmic resistance of the order of 10 S2 cm2. In more concentrated solutions the film is compact and in 0.5 M &Cl it may be only about 3 molecules thick. The aim of this paper was to use the PHMDE in chloride solutions for the study of the redox processes of some aromatic amines and to find the best conditions of film formation fos this purpose. EXPERIMENTAL
A hanging mercury drop electrode with a s&ace area-of (0.02 2 0.002) cm2 was used- as previously described [3] _ Chronovoftamperometric curves were ‘-recorded using a. Radiometer polarograph PO& ~411.solutions were deaerated wit$ electrolytically generated hydrogen. The supporting electrolyte was 0.1 bM %iClI cf. &&.lyt&al reagent grade. In several cases purifi* methanol wris add-ed ir?order~to improve the soltibility: of the-studibd compounds; All ‘organic substances -empioy.& were additionally purifiedby re&y-stallisation or dist$l~ .-_
-~ _
5
653
lation. A calomel electrode (0.1 M KCI) was used as reference. In some cases the charge~consumed in oxidation and reduction processes was measured with a microcoulometer @3] of a special type (polaroquanter)3 The passive film was always formed in the presence of the oxidizable species. RESULTS
The HMDE was polarized anodically in solutions of various concentrations of CI- ions and pH at scan rates 100, 200, 400 and 800 mV min-I. The optimum conditions for the subsequent anodic oxidation of the organic substances were found to be in solution of 0.1 M HCl at scan rate 800 mV min-‘. Measurements for solutions containing only the supporting electrolyte are shown as curve 1 on the following diagrams_ Peak (a) corresponds to the formation of calomel. The origin of peak (a’) is not well understood but it seems likely to be connected with some further reactions of mercury with Cl- ions. The presence of methyl alcohol- in solution had no significant influence on the possibility of oxidation, but as we have found previously 193 other anions such as NOT and ClO, act destructively on the passive layer and may not be tolerated above a concentration exceeding l/10 that of the Cl- ions.
of
Oxidation-reduction
benzidine
Results for experiments with benzidine are shown in Fig. 2. Peak (b) corresponds to the oxidation of benzidine. In the concentration range 2 X lo-’ -
+oJ3
.-
‘a4
E/y
Fig. 2. Ckrentholtage
tions of benzidine;
a’ curves
for 0.1
(1) 0, (2) 5 X.10m4
171 HCl solution_ containing the following concentraM, (3) l.0s3 M, (4) 2 5 10” M, (5) 4 X 10m3 M_
Fii. 3. Cyclic current-voltage
curves for 0.1 M HC‘I containing
10H3 M benzidine.
102M t.he height of peak (b) is directly proportional to the concentration of benzidine in the bulk solution. If the direction of the voltage sweep is reversed after the peak of benzidine oxidation is attained a corresponding cathodic peak appears (peak (c), curve 1, Fig. 3). The charges consumed in the benzidine oxidation and the subsequent reduction of the product formed were measured coulometrically with the polaroquanter and the ratio of charges was found to be 3: 1 when the cathodic curve was recorded immediately after registration of the oxidation peak. A multiple cyclic polarization under these conditions only slightly changes the heights of pe_ak (b) and (c) (curve 2, Fig_ 3): However, polarization of the e&&rode to more positive potentials makes the prccess highly irreversible
--Fig. 4. q>qent---volwe
CURVES for Q.l~M He1 sdlution containing (A) o-tohidine in co% cenGatiotir (I) 0, (2) 2 X 1UI3 Lx, (3) 4 _XI.0m3 M, (4) 8 X 10d3 ,?I, (.5) lO+ M,-(B) ptoluidine in cqncentrations: (I) 0, (2) LOW3 M, (3) 2 X 10H3 M, (4) 4 X 10V3 n/Ii (5) 3.0-a+ _
655
(curve 3, E’ig. 3) leading to. the compfete disappearance cf the cathodic peak (curve 4, Fig, 3). Thephenomena describedare similartothoseobsemed previously by us 14-j in the course -of oxidation of ferroeyanide on the PEIMDE, Oxidation of other amines Oxidation curves for isomeric o- and p-toluidines, p-aminoacetophenone, N-methylaniline, p-phenylenediamine and diphenyfamine are shown in Figs. 4-8, respectively. Oxidation of toluidines, p-aminoacetophenone and N-methylaniline occurs at rather positive potentials near +I V. In this potential region the electronic conductivity of the passive layer is poor and the oxidation currents were small. Furthermore, no &n-rents corresponding to reduction of the oxidation prodttcts formed could he observed. In the case of p-aminoacetophenone we found that the oxidation peaks were higher from solutions containing 10 ~01% of CHaOH than from aqueous solutions. Attempts to effect oxidation of m-toluidine were unsuccessful. Results for F-phenyfenediamine and diphenylamine were similar to those of benzidine. Initial oxidation of p-phenylenediamine occurs at a potential of
-+-+ I -i_ f
-z--
_!_
---I-+-I--k I -y--
+
L
7
5. Current-voltage curves for 0.1 fM HC1 + 10 vd% MeOH containing p-aminoacetophenone in concentrations: (I) 0, (2) lo-’ M, (3) 2 X low3 M, (4) 4 X low3 N, (5) 6 X 10v3 M, (6) 10--2&L Pig_
Fig. 6. Ckren~oltage curves for 0.1 -&IHCL containing N-methylaniline in concentrations: Q) 0, (2).5-X IO--~ N. (3) IO--? M, (4) 2 x 10-S M, (-5) 4 X 1O-3 AI, (6) 1O-2 I&
Fig. 7. Current-voltage tions: (1) 0, (2) 10e3 Fig- 8. Current-voltage in concentrations: (1)
M,
curves for 0.1 M HCl containing p-phenylenediarnine (3) 2 X lo-' M, (G%) 4 X 10M3 A.2.(5) 6 X 10e3 M_
in concentra-
curves for 0.1 .&I HCI -P-50 ~01% CH,OH containing diphenyiamine 0, (2) 2 x low4 kf, (3) 4 X 10B4 &f, (4) 6 X 10e4 1M, (5) low3 BE
+0.5 V at which the second peak in pure chloride solution also appears- Additional peaks at higher positive potentials shown in Fig. 7 indicate that the inG tial product undergoes further reactions. Oxidation of diphenylamine was studied in a solution containing 50 ~019%C;H30K, because of its low solubility. Again in the cases of both p-phenylenediamine and diphenylamme no reduction of the formed oxidation products co&d be observed. DfSCUSSION
The results showed that the Kemula type hanging mercury drop electrode [3], passivated in a solution of Cl- ions, can be a very convenient tool for studying the electrooxidation of organic compounds in the anodic potential rangeIn the studies by Kuwana and Adams Cl] and by Podbielancik [SJ of the oxidation of organic compounds, they used a passivated, mercury pool efectrade_-Podbielancik’s attempts to use the hanging mercury drop electrode of &he Gerischer type [lOI were unsuccessful- The-use of the mercury pool is a rather tedious technique because in order to obtain good reproducibility each experiment requires the-renewal of the solution and of the mercury. :Moreover in-both papers -[I J and [Zj it was-recommended that the organic substance should be added after the formation of the passive layer.
657
Our present experiments indicate similarly to the case of the ferrocyanide oxidation 141 that the presence of the oxidizable substances in solution before formation-of a passivating film has no visible influence on~the possibility of oxidation. It follows from our experiments that the possibility of organic electrooxidation largely depends on the structure and properties of the passive layer as well as on the kind of substituents and their position iy a molecule.
The most suitable experimental &angement for our purposes was the Hh!IDE covered with a rather thick and porous passive layer formed in dilute HCl solutions. In 0.1 M HC1 the aromatic amines are mainly in the form of cations with the ratio [RNHs]‘J[RNH,] > 103. The significance of the interaction of the -ir-electrons of the aromatic ring with the Hg surface for the process of adsorption of aromatic compounds at potentials corresponding to the positive branch of the electrocapillary curve has been shown by Frumkin [11] and subsequently by Gerovich and Rybal’ chenko LIZ]. In addition Blomgren and Bockris {13] as well as Conway and Bmdas [14] have stressed the significance of r-electron interaction with Hg in the adsorption process of various aromatic a-mines from acidic and neutral solutions. They have found that in acidic media the amine cations should lie flat on the Hg surface at potentials corresponding to the positive branch of the electrocapillary curve. We believe that when the mercury electrode is covered with calomel the interactions of rr-electrons of the aromatic ring with the passive layer should also play an important role in oxidation process. This view is supported by the fact that we were able to oxidize only aromatic amines. Attempts at oxidation of some aliphatic compounds were unsuccessful. If the oxidation reactions were diffusion controlled the plots of peak current vs- concentration should be linear with a slope in agreement with the RandlesdevCik equation [15] _ We only get such a dependence in the case of benzidine, where we assumed n = 2 and D = 5 X 10v cm2 s-l. The reversible oxidation of benzidine to diiminequinone was proved by Mat&a et aI- 11161 at TABLE
1
Comparisonof E1,s and Ep values at Pt and PHMDE,
respectively,
measured vs. 0.1 NCE
atpH=l
Compound
+JV
Eknzidine o-Toluidine PToluidine N-niethylaniline pAminoacetophenone Diphetiyiarnine p_Phenylenediamine
+0.580 +0.99 - io.93 +x.075 -4-l-02 +0.93 - +0.83 +1.05 +0.69 +0.52
%21V
0.57 0.82 0.79
[16] [lS) [IS]
_-.@jf3.
-. y-_
-_-
:
_- _-:
i 2
4
6
Fig. 9. Calibration Fig.
_ ._.
_.-
_- =
-
IO. Calibration
8
f0
10%/M curves
curves
for N-methylaniline for benzidine
(I),
(I),
p-toluidine
diphenylamine
(2)
and o-toluidine
(3).
(2) and p-phenylenediamine
(3).
the Pt disc electrode in a wide pH range (6-9):There is good agreement between the oxidation potentials at the Pt and PHMBE which suggest that the mechanism of benxidine oxidation is the same at both electrodesA comp_arison of other results is not relevant because the reactions are Meversible and-E, values depend on experimental conditions. Anodic oxidation of a series bf p-substituted anilines in aqueous solutions was studied at the carbon electrode by 3acon and Adams [17]_ They found a common mecha-nis_minvolving elimination of the para groups arid h&ad-to&ii coupling. The corresponding 4 substituted 4aminodiphenylamines were formed as oxidation products_ A sititiar mechanism may be postulated in the case- of the I p-substituted aniline-oxidation at *he PHMDE. Our experiments have.shown that only the ortl20and para-isomers could be oxid@d. Difficulties with the oxidation of_.~~eta:isomers are dtie to the lack of formation of the quinone strticturein this lattercase.The tialibration curves for the -compounds -which we have studied
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3 W.:_Kemulti&d 4 W: Ke-tiula, %
Z. tiLb!&, Anal; Chti. &&, 18 (l&,8) I&. Ktiblik and- d.~Taraszewska, J. EIectroanaJ. Chem., 6._(1963) 119. 5 *_ Ke@xuila,-2. J&blik-and J_ Taras&&vska, Proc:I~t. Symp_ on Microchemical Tech: ~n@ues, (19_Sl) 865: _ ----6 W. K&mula~and~J_ Taraszewske, Re;; Chim. Minerale, 5 (1968) 535. -7 B. Behr and J:~Taraszewska, J_ EJectroam& Chem., 19 (1968) 373. 8 -J-W: Strtijek and J. Zak, J_ Electroan& Chem., 46 (1973) 435. 9 -2. Kublik and J_ Terasze+&ka, BuII. Acad_ PoJon_ Sci_ Ser. Sci. Chim., 10 (1962) 515. 10 R. Gerischer, Z. Phys. Chem., 20 (1953) 302. 11 A-N. Frumkin; Ergebn. exakt. Naturwiss., 7 (1928) 256. 12 M.A. Gerovich and G.F. Rybal’chenko, Zh. Fiz. Khim., 32 (1958) 109. 13 E. Biomgren and J.O’M. Bockris, J. Phys. Chem., 63 (1959) 1475. 14 B.E:.Conway and R.G. Barradas, Electrochim. Acta, 5 (z961) 319, 349. 327; A SevEik, Collect. Czech. Chem. 15 J.E_B. Randles, Trans. Faraday Sot., 44 (1948) .Commun., 13 (1948) 349_ 16 M. Matrka, J. Pipaloti, Z. Sdgner and J. Marhold, ChemP&ml., 21 (1971) 14. 17 J_ Bacon and R.N. Adams, J. Amer. Chem. Sot., 90 (1968) 6596. 18 P-G. Desideri, L, Lepri and D. Heimler, J. Electroanal. Chem., 52 (1974) 93, 105.