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Metal oxide electrodes as sensors in complexometric titrations
rulantu. Vol. 23. pp. 503-508.Pergamon Press,1976.Printed in GreatBritain.
METAL OXIDE ELECTRODES AS SENSORS IN COMPLEXOMETRIC TITRATIONS ADAM HULANICKI
and
MAREK TROJANOWICZ
Institute of Fundamental Problems in Chemistry, University of Warsaw, Warsaw, Poland (Receioed 3 September 1975.
Accepted 12 January 1976)
Summary-EDTA titrations of lead and manganese, and of some other ions by using these as indicator ions, have been followed potentiometrically with PbO, and MnO, electrodes. Explanations are put forward for the anomalies observed in the titration curves. With the MnO, electrode, formation of an Mn(III)-EDIA complex is responsible for the diminished potential breaks. With the PbO, electrode the decrease in the potential break and the distortion of the curves in acid solution may be due to reaction between electrode coating and titrant. The MnO, electrode has been used for titrations of Ba*+, Ca2+, Cd2+ and Cu2+ with errors < 1%. Ion-selective membrane electrodes have tended to re.place metal and amalgam electrodes for end-point indication in potentiometric titrations. However, it seems that for several ions metal oxide electrodes may be useful either in direct potentiometry or as endpoint sensors in titrations. Oxide electrodes were originally used for the measurement of PH.’ The potential of an electrode consisting of a metal Me and its oxide MeO,,, may be expressed by E = Ei,.+,,
K1O - y
+ $ln
ln ao-
where KS0 represents the solubility product oxide. This equation can be transformed into
(1) of the
EXPERIMENTAL
When the metal oxide electrode is in contact with a solution of the metal ions in a different oxidation state from that in the oxide, then the potential-determining reaction may be described by MeO,,,
+ nH+ + (n - m) e$
Memi + nH,O
and the potential
given by
E = Eiro,l,Mcm+
RT - (n ___- m)F In aM,m+ n -f (n _ m)
tiometric titrations. Nomura and Nakagawa4 used a Pt/Mn02 electrode as indicator for the complexometric titration of various ions. Vandael’ used a platinum helix coated with lead or bismuth oxide for the same purpose. Wang and co-workers’ detected titration end-points with an electrode consisting of Pb02 in paraffin. The shapes of the titration curves were often unexplained. In this study the application of platinum electrodes, coated electrolytically with MnOz or PbOz, as indicator electrodes in chelatometric titrations has been investigated. A mechanism is put forward as an explanation of the observed irregularities in their behaviour.
Potentials were measured with Radiometer pHM 26 and Wroclaw N-512 ELPO pH-meters. Radiometer K 401 calome1 and G 222 B glass electrodes were used, and indicator electrodes were prepared by anodic deposition of oxides on a 0.7-cm2 platinum foil under conditions used in electrogravimetry? All reagents were of analytical grade. Doubly distilled water was used for the preparation of solutions. The concentrations of standard solutions were established by the usual methods. DISCUSSION AND RESULTS
RT k_ In
Direct chelatometric
titrations
The variation of the electrode potentials as a function of pH was studied for 10m3M solutions of manor, when the pH and the ionic strength are constant, ganese(I1) and lead(I1) with and without an excess of EDTA (Fig. 1). The curve 1 for Mn*+ is linear up RT to pH 9 with slope of 120 mV/pH, in good agreement (4) E = E$,=cons,. - (n ~ _ m)F In %m+ with the 118 mV/pH predicted by equation (3). The curvature at high pH is due to the formation of hydAccording to this equation the oxide electrode could. in principle, be used for a direct measurement roxo-complexes. Titrations should be possible above .of Me”‘, or in a potentiometric titration of this ion, pH 8, but it was found that stable potentials are or of other ions by using the technique of an indicator reached only very slowly, which make the titrations ion.’ rather impractical. The potentials of the PbO, electrode were much Polarized platinum electrodes coated with oxides of Mn(IV), Pb(IV) or Bi were used by Kraft3 in potenless reproducible than those of the Mn02 electrode 503 f
aH+
c3)
504
ADAM HULANICKIand MAREK TR~JANOWICZ
2
4
6
6
IO
2
4
6
6
PH
Fig. 1. Potential-pH relationship and Pb’+ Ions, respectively:
for MnO, l,l’--10W3
(u) and PbO, (hj electrodes in solutions containing Mn’+ M Me*+; 2,2’-10-3 M MeEDTA + 10m3 M EDTA.
and the agreement between the calculated and experimental values for a titration was poor (Fig. 2). The rate of titration significantly influences the results, which usually have a positive error (2-10x for titrations in acetate and ammonia buffers). In the first part of the titrations the potential of the Pb02 electrode is adequately described by the equation: RT
E=
GoLPh~’-+---
[Pb’+]
ln[H+]
+ y
10
PH
(5)
c(Ph
where apb is the side-reaction coefficient for Pb2+ taking into account the presence of hydroxide and buffer. However, at pH < 6.5 for a lead concentration of about IO-‘M the system PbOz, Pb2+ is thermodynamically unstable in aqueous solution. When’ the titrant is added in excess, the potential after the endpoint should be given by the equation: + ‘$ln[H’]
(6)
positive and the reaction 2Pb0,
+ 2H,Y2-
s 2Pb(II)Y’-
Fig. 2. Complexometric
I
Chelutometric
titrations with indicator ion
In this type of titration the potential of the indicator electrode’ should be given by the equafion: E = Eia2,~ez+ -
+ ‘qln[H’]
0
Fraction
titrated
titration
of 10-j M solutions
(7)
tends to proceed to the right. As a result of this the ratio [Y’]/[PbY’] decreases and the potential slowly decreases after the end-point. Wang found6 that the maximum sometimes observed for this titration curve does not correspond exactly with the end-point. Reaction (7) is a source of positive error even when the titration curve has a conventional S-shape. Because of these errors, the titration does not seem likely to have any useful applications.
where Y represents the EDTA ligand, and K,$,Y the conditional stability constant of the complex. However the potential of the system PbOz. PbY is much more
0
+ O2 + 2Hz0
I
Fraction
of Mn’+
PbO, (pH 4.6) indicator electrodes: solid lines-theoretical
titrated
and Pb2+ with MnO,
(pH 9.3) and curves; poinls--experimental data.
(8)
505
Metal oxide electrodes as sensors
Mn3+ + e$Mn”’
(11)
Mn3+ + Y”- $ Mn(III)Y-
(12)
Using the standard potentials, the following equation for the electrode potential can be arrived at: E = 2E;n02iMn~4 - E;,stiMnz+ + RT in
F
~Y4-12&wn,uCH+14
(13)
[Mn(III)Y -1
To compare this expression with experimental data the concentration of Mn(III)Y - should be eliminated, as it cannot be estimated directly. From the stability constant of the complex [Mn~III)Y-]
K Mn(iii)Y =
(14)
[Mn3+,[Y”-J
and the equilibrium constant L
1
a Fraction
I
I
I
2
titrated
(1.9
Fig. 3. Titration of 10-3M solutions of various metal ions with MnO, indicator electrode in 1M ammonia buffer at pH9 (Cd’+, CL?+) and at pH 10 (Ba’+,Ca”), CM, = 1o-5A4.
where N represents the ion titrated, CL the total concentration of the indicator ion and K’ the conditional stability constant of the complex indicated by subscript. In the case of the MnOz electrode, titrations of Ba’+, Ca’+, Cd’+ and Cu2+ with EDTA were investigated in various buffers, in the presence and absence of dissolved oxygen (Fig. 3). The ~tential breaks observed are adequate for a determination of the endpoint when concentrations of 10-3-10-4M are used. As expected, the concentration of the indicator ion (10-‘&f) does not affect the change in potential. The potentials observed at the beginning of the titrations (f = 0) are close to those predicted theoretically (Figs. 4,s) but those observed after the end-point (f = 2) are not, the difference being 150 or 190mV in the presence or absence of oxygen, respectively. independent of pH and the cation being titrated. Similar discrepancies have been observed by Nomura and Nakagawa.4 To explain this behaviour we have considered the following reaction: MnO, + Mn(II)Y’-
of the redox reaction MnOz + Mn2+ + 4Ht + 2Mn3’
+ 2H20
u B w
(16)
I a
-
0 E ew
ITIV
IO
200 SCE
+ H2Y2- + 2H’ -+ 2Mn(III)Y-
+ 2H,O
(9)
studied previously by Yoshino et ~1.~ Because of formation of Mn(III)Y- the titrated solutions become yellowish after the end-point and the disproportionation of Mn(III)Y- gives a dark precipitate of MnOz. In calculating the electrode potential the following equilibria should be considered: MnO, + 4H+ + 2e s Mn’+ + 2HZ0
(10)
E exp .
mV vs. SCE
Fig. 4. Comparison of theoretical and experimental potentials of MnO, indicator electrodes in titrations of various metal ions in 1M ammonia buffer in the absence of oxygen for C”Mn= lO-‘M atf= 0 andf= 2.
506
ADAM HULANICKIand MAREKTROJANOWICZ
200 -
Table 1. Determination of various metal ions with MnO, indicator electrode
f=O
Amount, pm& Error,
Metal
: $ %
I a
_u 8 w
- IO
0
200
Eexp*
Taken
Found
%
10.0
91.1
97.4 96.8 97.5
+0.3 -0.3 +04
Ca’ +
10.0
95.3
95.7 95.5 95.3 954 95.4
+0,2 -0.2 -01 -0.1
Cd’+
9.0
96.8
97.3 97.2 97.2 97.5
+@5 +0,4 fO.4 +0.7
cl?+
9.0
100.1
101.0 100.6 101.2 100.3
+0,9 +0,5 fl.1 + 0.2
ion
PH
Ba’+
mVvs. SCE
calculated for Cy = 10m3M, and using the standard potentials9 the value of AE is 177 mV, in very good agreement with the experimental values, thus lending support for the proposed reactions scheme. Though the formation of Mn(III)Y- is not a disadvantage in a practical titration, the slowness of the electrode to give stable potentials is. However, with a 3-min wait after addition of titrant, reasonable results were obtained (Table 1). With the PbOz electrode the curve of potential vs. pPb does not depend on the presence of other cations if no Pb’+ is added. From the potential us.
When
I
I
I
200
400 E exp .
mV vs. SCE
Fig. 5. Comparison of theoretical and experimental potentials of MnO, indicator electrodes in titrations of various metal ions in 1M ammonia buffer in the presence of 0 = 10-‘M atf= 0 andf= 2. oxygen for C,,,
calculated from the redox potentials as equal to 10-8.4”, one can obtain the following expression for the concentration of Mn(III)Y - : [Mn(III)Y-]
= KMn(,,,,YK1”[Mn2’]‘12Cy4-]
[H’]’ (17)
which may be substituted as:
in (13) to give the potential
+ R_TIn CY”-ICH’I” K”2[Mn2+]“2 F
(18)
When no Mn(III)Y- is formed, which is assumed to be the case, the potential of the electrode should be described by the equation: E = EinOllMnz+-
‘$ h
CH'I' [Mn2+] ‘I2
(19)
The difference between (19) and (18) is independent of both the cation titrated and the pH of the solution, and is given by AE = E~nOzmn~+ - E&J+~~z+
+ cln RT 1
G 1x
(20)
PH
Fig. 6. Potential-pH relationship_ for PbOz _ ^electrode ~in lo-‘M solutions containing Pb”, Ca”, Cu” or Zn”.
Metal
oxide
electrodes
as sensors
507
70c )w M k! z -
sot )-
E E 2 LY
30( )-
I
4
4
6I
1
I
8
I 10
I
I
I 12
PH
Fig. 7. Potential-pH relationship for complexometric titration indicator electrode: 1,1’-10-3 M Me’+ ; 2,2’-lO_sM
pH relationship (Fig. 6) it follows that in all instances the position of the equilibrium of the reaction Pb02
+ 2H+ z$ Pb’+ + 1/202
+ Hz0
Table
of Cu2+ (a) and Zn2+ (b) with PbOz MeEDTA + 10e3 M EDTA.
2. Determination of various metal indicator electrode
(21)
depends on the lead-ion concentration, and is shifted to the right to such an extent, that in titrations of other cations it is not necessary to add Pb2+ as indicator ion. As the potential-pH diagrams in the presence of Cu-EDTA and Zn-EDTA (Fig. 7) indicate that below pH 8, Pb02 should dissolve from the electrode, titrations should only be possible at pH higher than this, which was confirmed by experimental titration curves (Fig. 8). It was found, contrary to what was stated by Wang,6 that titrations in acetate medium were of no practical utility since the electrode reacts with the solution. Further, titrations in alkaline media are not recommended because of the large positive errors (Table 2).
Amount, Metal ion Ca’ +
PH 9.0
Taken
ions with PbOz
mtnole Found
0.0973
0.1047 0.1040 0.1050
+ 7.6 + 6.9 + I.9
10.1 Gl2+
9.0
0.1050
0.1065 0.1090
+1,4 +3,8
ZnZ+
8.0 9.1
0.0999
0.1035 0.1042 0.1025 0.1070 0.1040
f3.6 f4.3 + 26 f7.1 +4.1
9.4 9.8
Both the experiments and the calculations clearly indicate that the MnO, and PbO, electrodes cannot perform reliably as sensors in titrations, because they
(2 ) 56C ,
52C,> E
0
0 Fraction
titroted
Error, ”‘”
I
Fraction
titrated
Fig. 8. Complexometric titration with PbO, indicator electrode in 0.1 M acetate buffer (1.1’) and 0.1 M ammonia buffer (2.2’): 1,2-10-3 M Cuz+; 1’,2’-10-s M Cd*+.
ADAM HULANICKIand MAREK TROJANOWICZ
508
are thermodynamically conditions,
although
under certain electrodes
circumstances.
reported
only on the relative some conditions, tials should
unstable the MnO,
by other slowness
but stable
under electrode
the titration can be used
The application workers
of PbO,
must be based
of the reactions and reproducible
under poten-
not be expected.
REFERENCES 1. D. J. G. Ives and G. J. Janz, Reftirence Electrodes, Academic Press, New York, 1961.‘ 2. A. Hulanicki and M. Trojanowicz, Tulanta, 1969, 16, ,?C
3. G. Kraft, Z. Anal. Chem., 1962, 186, 187; 1969. 245, 58. 4. T. Nomura and G. Nakagawa, Bunseki Kaguku, 1967, 16, 1314. 5. C. Vandael, Ind. Chim. Be/ye, 1962, 8, 932. 6. C. N. Wang, P. J. Kinlen, D. A. Schoeller and C. 0. Huber, Anal. Chem., 1972, 44, 1152. .7. L. Meites, Handbook of Anulyticul Chemistry, McGraw-Hill, New York, 1963. 8. Y. Yoshino, A. Ouchi, Y. Tsunoda and M. Kajima, Can. J. Chem.. 1962. 40. 775. 9. K. J. Velter and G. Manecke. Z. Ploys. Chrrn. Leipzig, 1950. 195. 270.
METAL OXIDE ELECTRODES AS SENSORS IN COMPLEXOMETRIC TITRATIONS ADAM HULANICKI
and
MAREK TROJANOWICZ
Institute of Fundamental Problems in Chemistry, University of Warsaw, Warsaw, Poland (Receioed 3 September 1975.
Accepted 12 January 1976)
Summary-EDTA titrations of lead and manganese, and of some other ions by using these as indicator ions, have been followed potentiometrically with PbO, and MnO, electrodes. Explanations are put forward for the anomalies observed in the titration curves. With the MnO, electrode, formation of an Mn(III)-EDIA complex is responsible for the diminished potential breaks. With the PbO, electrode the decrease in the potential break and the distortion of the curves in acid solution may be due to reaction between electrode coating and titrant. The MnO, electrode has been used for titrations of Ba*+, Ca2+, Cd2+ and Cu2+ with errors < 1%. Ion-selective membrane electrodes have tended to re.place metal and amalgam electrodes for end-point indication in potentiometric titrations. However, it seems that for several ions metal oxide electrodes may be useful either in direct potentiometry or as endpoint sensors in titrations. Oxide electrodes were originally used for the measurement of PH.’ The potential of an electrode consisting of a metal Me and its oxide MeO,,, may be expressed by E = Ei,.+,,
K1O - y
+ $ln
ln ao-
where KS0 represents the solubility product oxide. This equation can be transformed into
(1) of the
EXPERIMENTAL
When the metal oxide electrode is in contact with a solution of the metal ions in a different oxidation state from that in the oxide, then the potential-determining reaction may be described by MeO,,,
+ nH+ + (n - m) e$
Memi + nH,O
and the potential
given by
E = Eiro,l,Mcm+
RT - (n ___- m)F In aM,m+ n -f (n _ m)
tiometric titrations. Nomura and Nakagawa4 used a Pt/Mn02 electrode as indicator for the complexometric titration of various ions. Vandael’ used a platinum helix coated with lead or bismuth oxide for the same purpose. Wang and co-workers’ detected titration end-points with an electrode consisting of Pb02 in paraffin. The shapes of the titration curves were often unexplained. In this study the application of platinum electrodes, coated electrolytically with MnOz or PbOz, as indicator electrodes in chelatometric titrations has been investigated. A mechanism is put forward as an explanation of the observed irregularities in their behaviour.
Potentials were measured with Radiometer pHM 26 and Wroclaw N-512 ELPO pH-meters. Radiometer K 401 calome1 and G 222 B glass electrodes were used, and indicator electrodes were prepared by anodic deposition of oxides on a 0.7-cm2 platinum foil under conditions used in electrogravimetry? All reagents were of analytical grade. Doubly distilled water was used for the preparation of solutions. The concentrations of standard solutions were established by the usual methods. DISCUSSION AND RESULTS
RT k_ In
Direct chelatometric
titrations
The variation of the electrode potentials as a function of pH was studied for 10m3M solutions of manor, when the pH and the ionic strength are constant, ganese(I1) and lead(I1) with and without an excess of EDTA (Fig. 1). The curve 1 for Mn*+ is linear up RT to pH 9 with slope of 120 mV/pH, in good agreement (4) E = E$,=cons,. - (n ~ _ m)F In %m+ with the 118 mV/pH predicted by equation (3). The curvature at high pH is due to the formation of hydAccording to this equation the oxide electrode could. in principle, be used for a direct measurement roxo-complexes. Titrations should be possible above .of Me”‘, or in a potentiometric titration of this ion, pH 8, but it was found that stable potentials are or of other ions by using the technique of an indicator reached only very slowly, which make the titrations ion.’ rather impractical. The potentials of the PbO, electrode were much Polarized platinum electrodes coated with oxides of Mn(IV), Pb(IV) or Bi were used by Kraft3 in potenless reproducible than those of the Mn02 electrode 503 f
aH+
c3)
504
ADAM HULANICKIand MAREK TR~JANOWICZ
2
4
6
6
IO
2
4
6
6
PH
Fig. 1. Potential-pH relationship and Pb’+ Ions, respectively:
for MnO, l,l’--10W3
(u) and PbO, (hj electrodes in solutions containing Mn’+ M Me*+; 2,2’-10-3 M MeEDTA + 10m3 M EDTA.
and the agreement between the calculated and experimental values for a titration was poor (Fig. 2). The rate of titration significantly influences the results, which usually have a positive error (2-10x for titrations in acetate and ammonia buffers). In the first part of the titrations the potential of the Pb02 electrode is adequately described by the equation: RT
E=
GoLPh~’-+---
[Pb’+]
ln[H+]
+ y
10
PH
(5)
c(Ph
where apb is the side-reaction coefficient for Pb2+ taking into account the presence of hydroxide and buffer. However, at pH < 6.5 for a lead concentration of about IO-‘M the system PbOz, Pb2+ is thermodynamically unstable in aqueous solution. When’ the titrant is added in excess, the potential after the endpoint should be given by the equation: + ‘$ln[H’]
(6)
positive and the reaction 2Pb0,
+ 2H,Y2-
s 2Pb(II)Y’-
Fig. 2. Complexometric
I
Chelutometric
titrations with indicator ion
In this type of titration the potential of the indicator electrode’ should be given by the equafion: E = Eia2,~ez+ -
+ ‘qln[H’]
0
Fraction
titrated
titration
of 10-j M solutions
(7)
tends to proceed to the right. As a result of this the ratio [Y’]/[PbY’] decreases and the potential slowly decreases after the end-point. Wang found6 that the maximum sometimes observed for this titration curve does not correspond exactly with the end-point. Reaction (7) is a source of positive error even when the titration curve has a conventional S-shape. Because of these errors, the titration does not seem likely to have any useful applications.
where Y represents the EDTA ligand, and K,$,Y the conditional stability constant of the complex. However the potential of the system PbOz. PbY is much more
0
+ O2 + 2Hz0
I
Fraction
of Mn’+
PbO, (pH 4.6) indicator electrodes: solid lines-theoretical
titrated
and Pb2+ with MnO,
(pH 9.3) and curves; poinls--experimental data.
(8)
505
Metal oxide electrodes as sensors
Mn3+ + e$Mn”’
(11)
Mn3+ + Y”- $ Mn(III)Y-
(12)
Using the standard potentials, the following equation for the electrode potential can be arrived at: E = 2E;n02iMn~4 - E;,stiMnz+ + RT in
F
~Y4-12&wn,uCH+14
(13)
[Mn(III)Y -1
To compare this expression with experimental data the concentration of Mn(III)Y - should be eliminated, as it cannot be estimated directly. From the stability constant of the complex [Mn~III)Y-]
K Mn(iii)Y =
(14)
[Mn3+,[Y”-J
and the equilibrium constant L
1
a Fraction
I
I
I
2
titrated
(1.9
Fig. 3. Titration of 10-3M solutions of various metal ions with MnO, indicator electrode in 1M ammonia buffer at pH9 (Cd’+, CL?+) and at pH 10 (Ba’+,Ca”), CM, = 1o-5A4.
where N represents the ion titrated, CL the total concentration of the indicator ion and K’ the conditional stability constant of the complex indicated by subscript. In the case of the MnOz electrode, titrations of Ba’+, Ca’+, Cd’+ and Cu2+ with EDTA were investigated in various buffers, in the presence and absence of dissolved oxygen (Fig. 3). The ~tential breaks observed are adequate for a determination of the endpoint when concentrations of 10-3-10-4M are used. As expected, the concentration of the indicator ion (10-‘&f) does not affect the change in potential. The potentials observed at the beginning of the titrations (f = 0) are close to those predicted theoretically (Figs. 4,s) but those observed after the end-point (f = 2) are not, the difference being 150 or 190mV in the presence or absence of oxygen, respectively. independent of pH and the cation being titrated. Similar discrepancies have been observed by Nomura and Nakagawa.4 To explain this behaviour we have considered the following reaction: MnO, + Mn(II)Y’-
of the redox reaction MnOz + Mn2+ + 4Ht + 2Mn3’
+ 2H20
u B w
(16)
I a
-
0 E ew
ITIV
IO
200 SCE
+ H2Y2- + 2H’ -+ 2Mn(III)Y-
+ 2H,O
(9)
studied previously by Yoshino et ~1.~ Because of formation of Mn(III)Y- the titrated solutions become yellowish after the end-point and the disproportionation of Mn(III)Y- gives a dark precipitate of MnOz. In calculating the electrode potential the following equilibria should be considered: MnO, + 4H+ + 2e s Mn’+ + 2HZ0
(10)
E exp .
mV vs. SCE
Fig. 4. Comparison of theoretical and experimental potentials of MnO, indicator electrodes in titrations of various metal ions in 1M ammonia buffer in the absence of oxygen for C”Mn= lO-‘M atf= 0 andf= 2.
506
ADAM HULANICKIand MAREKTROJANOWICZ
200 -
Table 1. Determination of various metal ions with MnO, indicator electrode
f=O
Amount, pm& Error,
Metal
: $ %
I a
_u 8 w
- IO
0
200
Eexp*
Taken
Found
%
10.0
91.1
97.4 96.8 97.5
+0.3 -0.3 +04
Ca’ +
10.0
95.3
95.7 95.5 95.3 954 95.4
+0,2 -0.2 -01 -0.1
Cd’+
9.0
96.8
97.3 97.2 97.2 97.5
+@5 +0,4 fO.4 +0.7
cl?+
9.0
100.1
101.0 100.6 101.2 100.3
+0,9 +0,5 fl.1 + 0.2
ion
PH
Ba’+
mVvs. SCE
calculated for Cy = 10m3M, and using the standard potentials9 the value of AE is 177 mV, in very good agreement with the experimental values, thus lending support for the proposed reactions scheme. Though the formation of Mn(III)Y- is not a disadvantage in a practical titration, the slowness of the electrode to give stable potentials is. However, with a 3-min wait after addition of titrant, reasonable results were obtained (Table 1). With the PbOz electrode the curve of potential vs. pPb does not depend on the presence of other cations if no Pb’+ is added. From the potential us.
When
I
I
I
200
400 E exp .
mV vs. SCE
Fig. 5. Comparison of theoretical and experimental potentials of MnO, indicator electrodes in titrations of various metal ions in 1M ammonia buffer in the presence of 0 = 10-‘M atf= 0 andf= 2. oxygen for C,,,
calculated from the redox potentials as equal to 10-8.4”, one can obtain the following expression for the concentration of Mn(III)Y - : [Mn(III)Y-]
= KMn(,,,,YK1”[Mn2’]‘12Cy4-]
[H’]’ (17)
which may be substituted as:
in (13) to give the potential
+ R_TIn CY”-ICH’I” K”2[Mn2+]“2 F
(18)
When no Mn(III)Y- is formed, which is assumed to be the case, the potential of the electrode should be described by the equation: E = EinOllMnz+-
‘$ h
CH'I' [Mn2+] ‘I2
(19)
The difference between (19) and (18) is independent of both the cation titrated and the pH of the solution, and is given by AE = E~nOzmn~+ - E&J+~~z+
+ cln RT 1
G 1x
(20)
PH
Fig. 6. Potential-pH relationship_ for PbOz _ ^electrode ~in lo-‘M solutions containing Pb”, Ca”, Cu” or Zn”.
Metal
oxide
electrodes
as sensors
507
70c )w M k! z -
sot )-
E E 2 LY
30( )-
I
4
4
6I
1
I
8
I 10
I
I
I 12
PH
Fig. 7. Potential-pH relationship for complexometric titration indicator electrode: 1,1’-10-3 M Me’+ ; 2,2’-lO_sM
pH relationship (Fig. 6) it follows that in all instances the position of the equilibrium of the reaction Pb02
+ 2H+ z$ Pb’+ + 1/202
+ Hz0
Table
of Cu2+ (a) and Zn2+ (b) with PbOz MeEDTA + 10e3 M EDTA.
2. Determination of various metal indicator electrode
(21)
depends on the lead-ion concentration, and is shifted to the right to such an extent, that in titrations of other cations it is not necessary to add Pb2+ as indicator ion. As the potential-pH diagrams in the presence of Cu-EDTA and Zn-EDTA (Fig. 7) indicate that below pH 8, Pb02 should dissolve from the electrode, titrations should only be possible at pH higher than this, which was confirmed by experimental titration curves (Fig. 8). It was found, contrary to what was stated by Wang,6 that titrations in acetate medium were of no practical utility since the electrode reacts with the solution. Further, titrations in alkaline media are not recommended because of the large positive errors (Table 2).
Amount, Metal ion Ca’ +
PH 9.0
Taken
ions with PbOz
mtnole Found
0.0973
0.1047 0.1040 0.1050
+ 7.6 + 6.9 + I.9
10.1 Gl2+
9.0
0.1050
0.1065 0.1090
+1,4 +3,8
ZnZ+
8.0 9.1
0.0999
0.1035 0.1042 0.1025 0.1070 0.1040
f3.6 f4.3 + 26 f7.1 +4.1
9.4 9.8
Both the experiments and the calculations clearly indicate that the MnO, and PbO, electrodes cannot perform reliably as sensors in titrations, because they
(2 ) 56C ,
52C,> E
0
0 Fraction
titroted
Error, ”‘”
I
Fraction
titrated
Fig. 8. Complexometric titration with PbO, indicator electrode in 0.1 M acetate buffer (1.1’) and 0.1 M ammonia buffer (2.2’): 1,2-10-3 M Cuz+; 1’,2’-10-s M Cd*+.
ADAM HULANICKIand MAREK TROJANOWICZ
508
are thermodynamically conditions,
although
under certain electrodes
circumstances.
reported
only on the relative some conditions, tials should
unstable the MnO,
by other slowness
but stable
under electrode
the titration can be used
The application workers
of PbO,
must be based
of the reactions and reproducible
under poten-
not be expected.
REFERENCES 1. D. J. G. Ives and G. J. Janz, Reftirence Electrodes, Academic Press, New York, 1961.‘ 2. A. Hulanicki and M. Trojanowicz, Tulanta, 1969, 16, ,?C
3. G. Kraft, Z. Anal. Chem., 1962, 186, 187; 1969. 245, 58. 4. T. Nomura and G. Nakagawa, Bunseki Kaguku, 1967, 16, 1314. 5. C. Vandael, Ind. Chim. Be/ye, 1962, 8, 932. 6. C. N. Wang, P. J. Kinlen, D. A. Schoeller and C. 0. Huber, Anal. Chem., 1972, 44, 1152. .7. L. Meites, Handbook of Anulyticul Chemistry, McGraw-Hill, New York, 1963. 8. Y. Yoshino, A. Ouchi, Y. Tsunoda and M. Kajima, Can. J. Chem.. 1962. 40. 775. 9. K. J. Velter and G. Manecke. Z. Ploys. Chrrn. Leipzig, 1950. 195. 270.