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Polarographic reduction of some anti bacterial compounds formed by the coupling of sulpha drugs with 8-hydroxy quinoline
WAHID; U. k&K Depa&ttetg (Rcccived7th
&d:P.ti.
DUA
of Chetn&ty,~ U&b&ity N&ember
4 &orkee.
1980; in r&vised foti
Raorkee 247672 (India)
25IhMay 1981)
ABSTRAC+ The polakgraphic reduction of coupled p&du&$ of diazonium salts of IO sulpha drugs with 1.3-dimethyl barbituric acid gave a single, tweelectron transfer, well-defined. diffusion-controlled, irreversible wave in Button-Robinkn.(BR) buffers of pH range 3.3-l 1.2; The mechanism of the reduction pre and the double-layer effects on the electrode~reactioa have been discussed.
INTRODUCTION
Like sulpha drugs, benzene sulphonamide, p-[(&hydroxy quinolyl) azo] and analogous compounds exhibit antibacterial activity [I]. The azo group of this compound is reduced mainly in liver and intestine [2] to form sulphanilamide and gihydroxy quinoline which inhibit the bacteria. The azo group is also polarographicaliy reducible. It Was -therefore thought worth while to reduce these medicinally important. compounds polarographi&lly at. t&e dropping mercury electrode (!WfE) and s&y. their hehaviour. comprehensively. .. F&ninary _Stud&s ‘revealed~~that they undergo a single two-electron transfer reduc& in & B&&-Robinson (BR) buffer over the. pH range 3.3- 11.2. A detailed accoumof studies‘of .-the mechanism -of the electrode &&ion are reported in .&is pap&.
212
structure: HO
R=H R=
compound (I) Benzene suIphonamide,p-[(S-hydroxyquinolyl)azo]
3
compound
(II)
R=
NH2
R=
compound(II1)
compound(V)
compound
(VI)
R=
compound (VII) =3
compound
R=
compound
(VIII)
R=
P
tJIN
compound (IX)
I
(X)
Reagents and solutions Stock solutions (1 X IOm-3M) of all these compounds were prepared in DMF(AR),. and BR buffers and various supporting electrolyte solutions were prepared in twice-distilled water. Twecn-20, Tritonex-35, Cetyltrimethyl ammonium iodide and sodium dodecyf benzene sdfonate were used to study the effect of surfactants.on the. polarographic reduction waves. Their solutions were prepared in twice-disti&d water. Apparatus The polarograms were recorded on a Cambridge pen recording polarograph. The capillary characteristics were 2.105 rngT13 s-‘/~ at h = 50 cm. The temperature was maintained at 30 f 0.2”C. Procedure The solution containing 1.0 ml of 1 X 10 -’ M depolariscr, 2.0 ml of DMF,. 1.0 ml of 1.0 M supporting electrolyte and 6.0 ml of buffer was analysed in a Lmgane
With $iogks&e &c~olysis the ykllow:brown &l&r of the catholyte started fad+& and a colourless soluti& waS o&ined after coniplete electiolysis. The latter f&t -we check& by the non-~ppearan&~ of. thy reductkjn way? ‘of the product of electrolysis and by the disappe+rance of- the tibsorption peak in the- visible region. The electrolysed solution- gave a posit&e. test for the hydrazo group [7] and a negative dye-test, thus confkming the. formation of hydrazo compound as the final reduction producL RESULTS
AND
DISCI+SION
All the compounds under invekigation-gave a single two-electron reduction wave over the entjre pH_.rangF~(3.3-11.2). -Qpical po@ro&ams are shown in Fig. 1, and the polarographic data are given in -Table 1. The li&ing current -of these waves was found to .be diffusion controlled, as skown frQm +e linear plot of ia vs. 6. The low
coinpoluld
--E&V
id/p&
..-
dE,/, dpH.-
: ?
:
pa-
’
Potential/V
-
-0.2V.curves(2)and Fig. 1. Typical pohrograms of azwzompounds (f)-(X) ( curves(l>and(6)startat (7) at -0.1 V. curves (3) and (4) at -0.3 V. curve (5) at -0.4 v). (1) ampound (I) at pH 6.2; (2) compound (II) at pH 4.2; (3) compound (IX) at pH 7.2; (4) compound (V) at pH 8.6; (5) compound (I) at pH 11.2:(6) compound(I) at pH 7.2: (7)compound(f) at pH6.2. Curves(l-5) and (7) at c= 1.0X IO-’ M, cure (6) at c=2.OX iO-‘ hf. Curves 1-6 at h=50 cm, curve 7 at h=60 cm.
value of the temperature .controlled
coefficient
of the current further confirmed
nature of the wave. The nature of the wave was found
the diffusion-
to be reversible
at
lo-’ M and below, over the entire pH range (3.3-11.2) but at higher concentration, viz. 1.O X 10 -’ M, it deviated from the reversible behaviour in the above pH range. Similar polarographic behaviour, viz reversibility at pH < 1 and pH > 13, but deviations from reversibility in the intermediate pH range for solutions of concentration higher than 5 X 10 + M, was observed [8-IO] in the case of azo-compounds. Such deviation has been ascribed to the adsorption of the depolariser [8]. The half-wave potential was found to be pH dependent up to 9.9 and shifted linearly towards more negative potential witb increase at pH (Fig. 2). The nature of the pH dependence of E,,, would indicate protonation of the depohuiser prior to electron uptake, but as only a single two-electron wave is obtained, the usual mechanistic path H’, e, e, H+ or H+, e, H+ , e is difficult to envisage. Therefore, the following mechanism for the electrode process can be put forward for the pH-dependent polarographic reduction: concentrations
5X
Ar-N=N-Ar+H+*Ar-GH=N-Ar Ar-
G H = N-Ar
+ 2 e + H+ P ArNH-NH-Ar
Hydrazo compound Au undoubtedly similar mechanism: Ar-N=N-Ar+2e+2H+#ArNH-NH-Ar
215
11
23L56789lQ
12
PH
Fig. 2. Plots of E,,, vs. pH for benzene sulphonamide, compound (I); 2: (III); 3: (IV); 4: (Ii).
p[(S-hydroxy-quinonlyl)azo]
series. Plot I:
can be proposed for the pH-independenttwo-electronwave, but thereare reasonsto believe that the electrode reaction follows the sequence e, IS+, e, H + and not e, e, H+, H+: Ar-N = N-Ar + e = [Ar-N = N+r]-
(A)
[Ar-N =N-Ar]=+e*[Ar-N=N-At]‘-
(B)
for-N
= N-A~]~-
+2 H+ + A~NH-NH-A~
.
kthe above -scheme, a monoanion radical (A) is formed by accepting one electron, followed by .formation--of, dianion (B). This should, however, be possible only in an .aprotic media [8]. But in protic medium, as used in the presentstudies, the basic monoanion radical can easily bedome protonated to give. the protonated kadikxl (C). This neutral free radical is iriunediately reduced by accepting one ele&ron apd proton to give a- single two-electron wave '._ : ‘[A~N,~N-A~]T+H+.~~* [-$I-NH-] '(c) [-X&NH-]
+ e Fsc [-N-NH-] . ..
- 5 A~NH-NH+& . ,Hydr azoknpound-
1.The_controlled-:pot+ia ekctrolysis resultqconfirm the formation of hydrazoc+mp&id and thereby’support this me&anism. This mechanism is- further.supported by-the wqrlc af oeers [9)..On ak-compouniis: ,.
Substituent effect In all electronegative groups studied, a shift of half-wave potential to more negative values result&. The effect of substituents increase in the sequence:
-C-CH
.
II]
q---J
’ I Cd%
Ordinarily, the electron-withdrawing groups should shift the E,,Z potential to a more positive value and a posiiive value of p, the specific reaction constant, is obtained. Since in our case E,,, shifts to more negative potentials (Table I), the adsorption effect appears to be more dominant. This is understandable in view of the presence of large adsorbable groups in these compounds. Effect of ionic strength Polarograms of compound I (concentration = 5 X 10 -’ M) were recorded using varying concentrations of KC1 at pH 4.2. As expected for a reversible system, no change in El,., or i, was observed. Effect of cations To determine the effect of the size and nature of the cation of the supporting electrolyte on the characteristics of the wave, the conipounds were polarographed using chlorides of lithium, sodium, potassium, tetramethyl and tetraethylammonium iodide. In all these supporting electrolytes the usual one well-defined wave was obtained. The El,* remained the same with lithium, sodium and potassium chloride, but shifted to more negative potentials with tetramethyl and tetraethyl halide (Table2). The latter behaviour can be attribuied to preferential adsorption of the tetramethyl and tetraethyl cation to that of the depolariser. This was further evident from the fact that the limiting current decreased only with these cations and not with others. Holleck et al. [ 111 reported that the adsorption lowered the concentratiqn of the depoladser on the electrode surface, as well as producing a different environment for the electron transfer by change of the double-layer structure. The combined effect of these two parameters would result in the lowering of the rate of the electrode process, subsequently shifting the. E,,, towards more negative potentials. with decrease in id. A further check was made by recording the polarograms with
217 T)?BLs
2
-:
.- ]
Effect of cations on Eln Support+g~ elc+olyte ..
.:
LiCl
and.&,of compound (V) at pH 8.6 ..
N&l
,_
KcI..
.. (CH,),N&
(C,H,).NI
sq2,/v
0.58
0.58
osa-
0.60
0.62
b/PA
0.425
0.425
0.425
0.4!3
0.40
supporting~eleckolytes having a common cation and different anions. As expixted, no change in I?,,, and i,- was observed, since the. half-wave potent&l of the compo*unds falls on the negative side-of the electrocapillary maxima. Effect of solvent cohpxition
The data on the variation of E,,, and i, wimincrease in the percentage of DMF, beyond. 30% (minimum, amount necessary to bring about tbc dissolution of the reactant) show that, with increase in percentage of DMF, the E,,i shifts almost linearly towards more negative potentials with simultaneous decrease in id (Table 3). An increase in the organic solvent content results in a rise in pH [12,13] and in an increase in the dissociation constant of the protonated species [14]. Both these factors lower the rate of protonation and consequently would lead to a shift in E,,, of the reduction wave towards more negative potentials in all such cases where protonation precedes the electron transfer. It appears that these two factors are not the only ones responsible for the observed shift in El,* in our case because the shift observed is greater than would be expected from the .change in pH and the dissociation constant (an increase in percentage of DMF from 30% to 60% resulted in an increase of pH by 1.0 unit). -This additional shift in A?&,~may be ascribed to a decrease in adsorbability and hence surface concentration of the depolariser with an increasing percentage of the
TABLE3
Effect of DMF conceitration on E ,/z and i, of compound (I) at pI!J 8.6 and compound (1II)at pH 7.2. c=IXIO-‘M Compound(I) % DMF -+2/V ihA
30
40
50
60
70 Ilkiermed wave
0.55
0.60
0:65
0.71
0.438
0.40
0.288
0.188
jo
40
0.47 O-425
0.52 0.394
Compound (III j
%DW. --E,,,/V
b/PA
50
.0.57 0.319
60
70
0.62 0.206
IR_dermed wave
218
in an aqueous-organic mixture [ 151. Obviously, decrease in surface concentration would retard the electrode process resulting.in a decrease-in QZ and &. As further assurance that the observed shift in El/z was not due merely to the changes in liquid-junction potential in the presence of water&DMF mixtures, polarograms using Rb+ as the piiot ion with the solvent composition (1 ml soln. of 0.05 M RbCl + 1 ml of 0.1 A4 (CH,),NBr + 3 ml DMF + 5 ml water) were recorded using a mercury pool as reference electrode_ The increase in percentage of DMF from 30% to 60% did not show any shift in E,,(E,, = -2_17V). However, 01% replacing Rb+ with azo-compound, the E,,, varied considerably with variation in the DMF content of. the solution (about -0.05 V per 10% rise in DMF). Since i, is dependent on the viscosity of the solvent, it was considered desirable to determine the viscosity of the solutions subjected to-electrolysis. The value of-i,,& decreases with amount of DMF added and this may be attributed to the changes in the solvation number (since no maximum is observed in the solvent compositionviscosity curves, the existence of association is ruled out)_ It can be argued that the increase in viscosity with the increase in DMF is not fully counterbalanced by the decrease in i,, (which is quite large due. to solvation of the depolariser), and hence the expected constancy id-/i; is not real&d. DMF
Effecr of swfactants
During the course of polarographic reduction either the electroactive species or the reduction product is normally adsorbed at the DME. The surface-active species, if any, is also preferentially adsorbed in a majority of cases. According to HoUeck [16], electrochemical reduction in the presence of surface-active admixtures can follow at least two reaction paths, viz.. “boundary surface reaction’? and ‘.‘solution reaction”. The first, which progresses more rapidly, can be excluded by covering of the boundary surface by surfactant molecules, thereby suppressing the adsorption of the depolariser. The rate of the electrode process is thus slowed down. However, in cases where the acceleration of the reduction process takes place, it would involve promotion of the boundary surface reaction in which a more rapid transfer of electrons to the proton&d molecule takes place. Such an effect was observed by Holleck [16] by the addition of gelatin in the reduction of azocompounds, and was explained in terms of proton conduction mechanism over the hydrogen bridges of the gelatin. Addition of surfactant in the EMC range shifts the E,,,z towards more negative potentials with simultaneous decrease in the limiting current. These observations can be explained in terms of slowness of the boundary surface reaction as propos& by Holleck. REFERENCES
1 L. Mod and A. Simoneni,Gazz. Ctim. Ital.. 70 (1940) 369. 2 R. Walker. Food Cosmet.Toxicol..8 (1970)659.
219 : 3 C, Com&c-.&I
Hi Arce,
C&n.’ Abstr; 49 (i955)
3251:
:-
-4 T. De&+x -&td J.L Kman, J. Ani Cbem. !ioc., .75 ( 1953) 2&l. .5 V_.Ne+dX~; Col!ect_ Cz+ C&m.Co~tm-; I (1922) 3196 J. Liigiie. tiG_ Swain aud;M. Fields, J. Am. C&x& Sot., 65 (!Q=+,
1349. : 7 .E: F&l. Spot:Te&s in‘O$gat& Aaalysis, E$+r. Neal York, 1966. p- 283. 8 -A.J_.&si-d&d H. Lund, Enjclopedia of Electrochemistry of_the Elememts. .Vol. ~111. Marcei Decker.
N&v York, 1973. p. 171. 9 V.B.~Ny@d, At+ Kern...28(I!+) 167. 10 I.M. Isa, R.M. Is& Y.M. Temerek and M.R.
M&,mqud, Ekctrocbim. Acta. 18 (1973) 139.
Ii B. Kastening, I-LGartmann aud L. Hokk, 12 K; !Mxw-Ad&ccs 13 K. schwabe, Fr@css
Ekctrkhim. Act& 9(6) (1964) 741. in Polarogmphy, VQI. III, Peqamon Press, London,. 1960, p. 911.
in Polaiogmphy, VoL I, Interscicncc, New York, 1462, pi 333.
14 S. +faiiauovskii ind V. Gultjaj, Ekctrokhimiya, I .(1%5)
76.
15 S Ma@novskii, Ta@nta, 12 (1%5) 1299. 16 G. Holkck and L. Ho&c&, Naturwissenschatten, 51 (1964) 212.
&d:P.ti.
DUA
of Chetn&ty,~ U&b&ity N&ember
4 &orkee.
1980; in r&vised foti
Raorkee 247672 (India)
25IhMay 1981)
ABSTRAC+ The polakgraphic reduction of coupled p&du&$ of diazonium salts of IO sulpha drugs with 1.3-dimethyl barbituric acid gave a single, tweelectron transfer, well-defined. diffusion-controlled, irreversible wave in Button-Robinkn.(BR) buffers of pH range 3.3-l 1.2; The mechanism of the reduction pre and the double-layer effects on the electrode~reactioa have been discussed.
INTRODUCTION
Like sulpha drugs, benzene sulphonamide, p-[(&hydroxy quinolyl) azo] and analogous compounds exhibit antibacterial activity [I]. The azo group of this compound is reduced mainly in liver and intestine [2] to form sulphanilamide and gihydroxy quinoline which inhibit the bacteria. The azo group is also polarographicaliy reducible. It Was -therefore thought worth while to reduce these medicinally important. compounds polarographi&lly at. t&e dropping mercury electrode (!WfE) and s&y. their hehaviour. comprehensively. .. F&ninary _Stud&s ‘revealed~~that they undergo a single two-electron transfer reduc& in & B&&-Robinson (BR) buffer over the. pH range 3.3- 11.2. A detailed accoumof studies‘of .-the mechanism -of the electrode &&ion are reported in .&is pap&.
212
structure: HO
R=H R=
compound (I) Benzene suIphonamide,p-[(S-hydroxyquinolyl)azo]
3
compound
(II)
R=
NH2
R=
compound(II1)
compound(V)
compound
(VI)
R=
compound (VII) =3
compound
R=
compound
(VIII)
R=
P
tJIN
compound (IX)
I
(X)
Reagents and solutions Stock solutions (1 X IOm-3M) of all these compounds were prepared in DMF(AR),. and BR buffers and various supporting electrolyte solutions were prepared in twice-distilled water. Twecn-20, Tritonex-35, Cetyltrimethyl ammonium iodide and sodium dodecyf benzene sdfonate were used to study the effect of surfactants.on the. polarographic reduction waves. Their solutions were prepared in twice-disti&d water. Apparatus The polarograms were recorded on a Cambridge pen recording polarograph. The capillary characteristics were 2.105 rngT13 s-‘/~ at h = 50 cm. The temperature was maintained at 30 f 0.2”C. Procedure The solution containing 1.0 ml of 1 X 10 -’ M depolariscr, 2.0 ml of DMF,. 1.0 ml of 1.0 M supporting electrolyte and 6.0 ml of buffer was analysed in a Lmgane
With $iogks&e &c~olysis the ykllow:brown &l&r of the catholyte started fad+& and a colourless soluti& waS o&ined after coniplete electiolysis. The latter f&t -we check& by the non-~ppearan&~ of. thy reductkjn way? ‘of the product of electrolysis and by the disappe+rance of- the tibsorption peak in the- visible region. The electrolysed solution- gave a posit&e. test for the hydrazo group [7] and a negative dye-test, thus confkming the. formation of hydrazo compound as the final reduction producL RESULTS
AND
DISCI+SION
All the compounds under invekigation-gave a single two-electron reduction wave over the entjre pH_.rangF~(3.3-11.2). -Qpical po@ro&ams are shown in Fig. 1, and the polarographic data are given in -Table 1. The li&ing current -of these waves was found to .be diffusion controlled, as skown frQm +e linear plot of ia vs. 6. The low
coinpoluld
--E&V
id/p&
..-
dE,/, dpH.-
: ?
:
pa-
’
Potential/V
-
-0.2V.curves(2)and Fig. 1. Typical pohrograms of azwzompounds (f)-(X) ( curves(l>and(6)startat (7) at -0.1 V. curves (3) and (4) at -0.3 V. curve (5) at -0.4 v). (1) ampound (I) at pH 6.2; (2) compound (II) at pH 4.2; (3) compound (IX) at pH 7.2; (4) compound (V) at pH 8.6; (5) compound (I) at pH 11.2:(6) compound(I) at pH 7.2: (7)compound(f) at pH6.2. Curves(l-5) and (7) at c= 1.0X IO-’ M, cure (6) at c=2.OX iO-‘ hf. Curves 1-6 at h=50 cm, curve 7 at h=60 cm.
value of the temperature .controlled
coefficient
of the current further confirmed
nature of the wave. The nature of the wave was found
the diffusion-
to be reversible
at
lo-’ M and below, over the entire pH range (3.3-11.2) but at higher concentration, viz. 1.O X 10 -’ M, it deviated from the reversible behaviour in the above pH range. Similar polarographic behaviour, viz reversibility at pH < 1 and pH > 13, but deviations from reversibility in the intermediate pH range for solutions of concentration higher than 5 X 10 + M, was observed [8-IO] in the case of azo-compounds. Such deviation has been ascribed to the adsorption of the depolariser [8]. The half-wave potential was found to be pH dependent up to 9.9 and shifted linearly towards more negative potential witb increase at pH (Fig. 2). The nature of the pH dependence of E,,, would indicate protonation of the depohuiser prior to electron uptake, but as only a single two-electron wave is obtained, the usual mechanistic path H’, e, e, H+ or H+, e, H+ , e is difficult to envisage. Therefore, the following mechanism for the electrode process can be put forward for the pH-dependent polarographic reduction: concentrations
5X
Ar-N=N-Ar+H+*Ar-GH=N-Ar Ar-
G H = N-Ar
+ 2 e + H+ P ArNH-NH-Ar
Hydrazo compound Au undoubtedly similar mechanism: Ar-N=N-Ar+2e+2H+#ArNH-NH-Ar
215
11
23L56789lQ
12
PH
Fig. 2. Plots of E,,, vs. pH for benzene sulphonamide, compound (I); 2: (III); 3: (IV); 4: (Ii).
p[(S-hydroxy-quinonlyl)azo]
series. Plot I:
can be proposed for the pH-independenttwo-electronwave, but thereare reasonsto believe that the electrode reaction follows the sequence e, IS+, e, H + and not e, e, H+, H+: Ar-N = N-Ar + e = [Ar-N = N+r]-
(A)
[Ar-N =N-Ar]=+e*[Ar-N=N-At]‘-
(B)
for-N
= N-A~]~-
+2 H+ + A~NH-NH-A~
.
kthe above -scheme, a monoanion radical (A) is formed by accepting one electron, followed by .formation--of, dianion (B). This should, however, be possible only in an .aprotic media [8]. But in protic medium, as used in the presentstudies, the basic monoanion radical can easily bedome protonated to give. the protonated kadikxl (C). This neutral free radical is iriunediately reduced by accepting one ele&ron apd proton to give a- single two-electron wave '._ : ‘[A~N,~N-A~]T+H+.~~* [-$I-NH-] '(c) [-X&NH-]
+ e Fsc [-N-NH-] . ..
- 5 A~NH-NH+& . ,Hydr azoknpound-
1.The_controlled-:pot+ia ekctrolysis resultqconfirm the formation of hydrazoc+mp&id and thereby’support this me&anism. This mechanism is- further.supported by-the wqrlc af oeers [9)..On ak-compouniis: ,.
Substituent effect In all electronegative groups studied, a shift of half-wave potential to more negative values result&. The effect of substituents increase in the sequence:
-C-CH
.
II]
q---J
’ I Cd%
Ordinarily, the electron-withdrawing groups should shift the E,,Z potential to a more positive value and a posiiive value of p, the specific reaction constant, is obtained. Since in our case E,,, shifts to more negative potentials (Table I), the adsorption effect appears to be more dominant. This is understandable in view of the presence of large adsorbable groups in these compounds. Effect of ionic strength Polarograms of compound I (concentration = 5 X 10 -’ M) were recorded using varying concentrations of KC1 at pH 4.2. As expected for a reversible system, no change in El,., or i, was observed. Effect of cations To determine the effect of the size and nature of the cation of the supporting electrolyte on the characteristics of the wave, the conipounds were polarographed using chlorides of lithium, sodium, potassium, tetramethyl and tetraethylammonium iodide. In all these supporting electrolytes the usual one well-defined wave was obtained. The El,* remained the same with lithium, sodium and potassium chloride, but shifted to more negative potentials with tetramethyl and tetraethyl halide (Table2). The latter behaviour can be attribuied to preferential adsorption of the tetramethyl and tetraethyl cation to that of the depolariser. This was further evident from the fact that the limiting current decreased only with these cations and not with others. Holleck et al. [ 111 reported that the adsorption lowered the concentratiqn of the depoladser on the electrode surface, as well as producing a different environment for the electron transfer by change of the double-layer structure. The combined effect of these two parameters would result in the lowering of the rate of the electrode process, subsequently shifting the. E,,, towards more negative potentials. with decrease in id. A further check was made by recording the polarograms with
217 T)?BLs
2
-:
.- ]
Effect of cations on Eln Support+g~ elc+olyte ..
.:
LiCl
and.&,of compound (V) at pH 8.6 ..
N&l
,_
KcI..
.. (CH,),N&
(C,H,).NI
sq2,/v
0.58
0.58
osa-
0.60
0.62
b/PA
0.425
0.425
0.425
0.4!3
0.40
supporting~eleckolytes having a common cation and different anions. As expixted, no change in I?,,, and i,- was observed, since the. half-wave potent&l of the compo*unds falls on the negative side-of the electrocapillary maxima. Effect of solvent cohpxition
The data on the variation of E,,, and i, wimincrease in the percentage of DMF, beyond. 30% (minimum, amount necessary to bring about tbc dissolution of the reactant) show that, with increase in percentage of DMF, the E,,i shifts almost linearly towards more negative potentials with simultaneous decrease in id (Table 3). An increase in the organic solvent content results in a rise in pH [12,13] and in an increase in the dissociation constant of the protonated species [14]. Both these factors lower the rate of protonation and consequently would lead to a shift in E,,, of the reduction wave towards more negative potentials in all such cases where protonation precedes the electron transfer. It appears that these two factors are not the only ones responsible for the observed shift in El,* in our case because the shift observed is greater than would be expected from the .change in pH and the dissociation constant (an increase in percentage of DMF from 30% to 60% resulted in an increase of pH by 1.0 unit). -This additional shift in A?&,~may be ascribed to a decrease in adsorbability and hence surface concentration of the depolariser with an increasing percentage of the
TABLE3
Effect of DMF conceitration on E ,/z and i, of compound (I) at pI!J 8.6 and compound (1II)at pH 7.2. c=IXIO-‘M Compound(I) % DMF -+2/V ihA
30
40
50
60
70 Ilkiermed wave
0.55
0.60
0:65
0.71
0.438
0.40
0.288
0.188
jo
40
0.47 O-425
0.52 0.394
Compound (III j
%DW. --E,,,/V
b/PA
50
.0.57 0.319
60
70
0.62 0.206
IR_dermed wave
218
in an aqueous-organic mixture [ 151. Obviously, decrease in surface concentration would retard the electrode process resulting.in a decrease-in QZ and &. As further assurance that the observed shift in El/z was not due merely to the changes in liquid-junction potential in the presence of water&DMF mixtures, polarograms using Rb+ as the piiot ion with the solvent composition (1 ml soln. of 0.05 M RbCl + 1 ml of 0.1 A4 (CH,),NBr + 3 ml DMF + 5 ml water) were recorded using a mercury pool as reference electrode_ The increase in percentage of DMF from 30% to 60% did not show any shift in E,,(E,, = -2_17V). However, 01% replacing Rb+ with azo-compound, the E,,, varied considerably with variation in the DMF content of. the solution (about -0.05 V per 10% rise in DMF). Since i, is dependent on the viscosity of the solvent, it was considered desirable to determine the viscosity of the solutions subjected to-electrolysis. The value of-i,,& decreases with amount of DMF added and this may be attributed to the changes in the solvation number (since no maximum is observed in the solvent compositionviscosity curves, the existence of association is ruled out)_ It can be argued that the increase in viscosity with the increase in DMF is not fully counterbalanced by the decrease in i,, (which is quite large due. to solvation of the depolariser), and hence the expected constancy id-/i; is not real&d. DMF
Effecr of swfactants
During the course of polarographic reduction either the electroactive species or the reduction product is normally adsorbed at the DME. The surface-active species, if any, is also preferentially adsorbed in a majority of cases. According to HoUeck [16], electrochemical reduction in the presence of surface-active admixtures can follow at least two reaction paths, viz.. “boundary surface reaction’? and ‘.‘solution reaction”. The first, which progresses more rapidly, can be excluded by covering of the boundary surface by surfactant molecules, thereby suppressing the adsorption of the depolariser. The rate of the electrode process is thus slowed down. However, in cases where the acceleration of the reduction process takes place, it would involve promotion of the boundary surface reaction in which a more rapid transfer of electrons to the proton&d molecule takes place. Such an effect was observed by Holleck [16] by the addition of gelatin in the reduction of azocompounds, and was explained in terms of proton conduction mechanism over the hydrogen bridges of the gelatin. Addition of surfactant in the EMC range shifts the E,,,z towards more negative potentials with simultaneous decrease in the limiting current. These observations can be explained in terms of slowness of the boundary surface reaction as propos& by Holleck. REFERENCES
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