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Potentiometric determination of hydrogen peroxide with permanganate in presence of fluoride, with iron as a catalyst: Application to the titanium—fluoride—peroxide mixed complex
@X9-9140/81 I20951-04$02.00/O Pergamon Press Ltd
r‘,l‘,“r,,. Vol. 2x. pp. 951 I0 954. 1981 Prmted I” Great Br~tam
SHORT
COMMUNICATIONS
POTENTIOMETRIC DETERMINATION OF HYDROGEN PEROXIDE WITH PERMANGANATE IN PRESENCE OF FLUORIDE, WITH IRON AS A CATALYST: APPLICATION TO THE TITANIUM-FLUORIDE-PEROXIDE MIXED COMPLEX L. POLO DIEZ, J. HERNANDEZ MENDEZ and M. J. ALMENDRAL PARRA Department
of Analytical
(Rrcrired
Chemistry.
6 Norrmhrr
Faculty
1980 Reaisrd
of Chemistry.
University
30 April 1981. Accrpred
of Salamanca
Spain
12 May 1981)
Summary-A potentiometric permanganate titration has been developed for the determination of hydrogen peroxide in the presence of fluoride, as well as of the peroxide content in the !itanium-Ruorideperoxide mixed complex. It is based on the stabilization of manganese(II1) with an excess of fluoride in a moderately acidic medium (pH close to 3) and on the use of iron(lIl) as catalyst. Errors are less than 0.5” 0
Several titrimetric methods are available for determination of hydrogen peroxide or the peroxide content of compounds, at macro-levels, those with cerium(IV) or permanganate being the most suitable.’ In the cerium(lV) methods, the presence of fluoride seems to lead to low results,’ but there is no information about its effect on the permanganate methods, although several permangate titrations in fluoride medium have been developed.3m9 In this paper, a potentiometric method is proposed for the determination of hydrogen peroxide with permanganate in the presence of the fluoride, the permanganate being reduced to manganese(M) stabilized as the fluoride complex. The method is applied to the determination of peroxide in the titaniumPfluorideperoxide mixed complex.
the potentials
between
a platinum
wire indicator electrode
and a saturated calomel electrode. Procrdurr Weigh accurately
about
0.1-0.5
g of the titanium-fluori-
de-peroxide complex (or use an equivalent amount of a solution containing this compound or hydrogen peroxide), dissolve it in a small volume of distilled water, and add 10 g of sodium fluoride. Add sufficient 1M sulphuric acid to give a pH close to 3 (Methyl Orange as indicator), enough Fe(lll) solution to give a final concentration about 10e4M. Titrate potentiometrically with the
and of per-
man@nate sohtion. RESULTS AND DISCUSSION Prdiminary
wsults
When the classic permanganate titration in acidic medium was used to determine the peroxide content in the complex, low but reproducible results were obtained. EXPERIMENTAL This error was attributed to partial stabilization of manganese as the manganese(II1) fluoride complex, which was confirmed by the pink colour observed Potassium permanganate solutions. about O.lN. i.e., before the characteristic permanganate colour 0.02M. were standardized against sodium oxalate. Hydroappeared. This fluoride interference cannot be gen peroxide solutions. about O.lN. were standardized avoided by heating in an acidic medium, as this against the permanganate. Other solutions included 1M i brings about decomposition of the peroxide group. and O.lM sulphuric acid. ISM sodium fluoride. and O.OlOM copper sulphate. chromic sulDhate. ammonium To overcome this problem the permanganate titra. metabanadate and ferric sulphate. tion was done in the presence of an excess of fluoride. The best conditions were determined by titrating hyS1o11plr Solid (NH,),Ti02F, was prepared by adding an excess drogen peroxide solutions. The characteristic pink of hydrogen peroxide and ammonium fluoride to a titaof the manganese(IIItfluoride colour complex nium(IV) sulphate solution. which was kept cold and neunecessitated the use of a physical method for the detralized with an ammonia solution to pH 7. termination of the end-point, the potentiometric method with a platinum wire as the indicator elecA Metrohm E-510 potentiometer was used to measure trode being the most suitable.
952
SHORT COMMUNICATIONS
900
-
700-
500-
300 -
100 -
1
t
L
75
loo
I
125
1
150
H*02 titrated, % Fig. I. Efiect of pH on the potentiometric curves. H202: 1.012 meq: NaF: l.OM; Fe(Il1): KMnO,: 0.026SM. pH values are those specifed on the curves.
Porc~llriot?lc~fr;~~
titrcttion
The reaction between permanganate and hydrogen peroxide in the presence of an excess of fluoride at about pH 3 is very slbw. especially near the equivalence point. Moreover. the electrode potential also stabilizes slowly. However. in the presence of copper. vanadium or iron salts. the chemical reaction occurs quickly. as shown by the instantaneous decolorization of permanganate. and the electrode potential stabilizes rapidly. Iron salts produced the best results and were therefore used as the catalyst. According to the literature. iron(III) peroxide species must be involved in the mechanism of the catalysed chemical reaction.1° The potentiometric titration curves obtained in the presence of iron(II1) were well defined, making determination of the end-point easy. A single titration took about IO min. normal for this kind of titration. The potentials obtained with excess of permanganate indicated that the electrode probably responded to a mixed potential. Efl?rr ofpH. Titration curves obtained in the presence of IO- ‘lLf iron(II1) at different pH values are shown in Fig. I. The change in potential around the equivalence point decreases when the pH is increased. and above pH 4 manganese dioxide is formed. The apparent tinal oxidation states of manganese (assum-
ing a quantitative reaction between permanganate and hydrogen peroxide) calculated from the titration curves of Fig. 1 are shown in Table I. These results lead to the conclusion that permanganate is not reduced stoichiometrically to manganese(II) even at pH I. On the other hand, at pH values close to 3. it is reduced quantitatively to manganese(“I). Taking into account that under these conditions the MnF4 complex must predominate in solution, the following reaction presumably occurs MnO;
+ 2H202
+ 4F-
+ 4H+ ---t MnF; + 20,
Titration error assuming the half-reaction MnO; + Mn(II1)
+ 4H,O
As is clear from the table,‘at pH values either lower or higher than 3. other reactions occur simultaneously with the formation of manganese(H) or (IV) respectively. The adjustment of pH to about 3 is easy owing to the presence of the HF/Fbuffer; some Methyl Orange indicator is added, followed by sufficient IM sulphuric acid to turn the colour red. A pH-meter is not needed and a glass electrode should not be used in such a fluoride medium anyway. Ejkt @the iron(lll) concmtration. Concentrations of Fe(II1) in the range 5 x 10-5P10-4M are suitable to make both the chemical and electrochemical reac-
Table I. Apparent tinal oxidation state of manganese and titration error. assuming the MnOJ -+ Mn(IIl) half-reaction at diRerent pH values PH Apparent final oxidation states of Mn
10m4M;
1.0
2.0
3.0
4.0
5.0
2.51
2.96
2.99
3.04
3.77
- 11.0
- I.1
LO.31
+ I.0
+20.3
SHORT
I
I
75
I 150
I 125
I 100
HaOa titrated, %
Fig. 2. EtTect of the NaF concentration on the potentiometric titration curves. Hz02: 0.949 meq: Fe(II1): 10e4M; pH: 3.0: KMn04: 0.0211M. NaF concentration (M); (1) 0.6: (2) 0.8: (3) 1.2 (4) 1.4. tions fast enough. However, as the titration times increased significantly for concentrations lower than 5 x 10-5M, 10e4M was taken as the optimum concentration. Eficr of sodium Juoride concentration. Titration curves obtained at different sodium fluoride concentrations are shown in Fig. 2. The change of potential around the equivalence point decreases slightly for increasing concentrations of sodium fluoride. However. the potential of the MnO;/Mn(lII) system should increase with the sodium fluoride concentration. The cause of the effect observed is obscure, but may be due to the effect of fluoride on the titanium complex as a whole. In other words, the increase in fluoride concentration will increase the concentration of the mixed complex and decrease the
900
953
COMMUNICATIONS
concentration of free peroxide in equilibrium with it, hence increasing the potential of the oxygen/mixedligand complex couple, and accounting for the higher potentials before the end-point. The lower potentials after the end-point could be due to a mixed-potential electrode response. l@ct of temperature. Titration curves obtained at different temperatures are shown in Fig. 3. At room temperatures in the range 15-25” no significant differences are observed in the shape of the titration curves. However, at temperatures higher than about 30”, changes in the titration curves are clear, and low results are obtained. This is probably due to disproportionation of hydrogen peroxide at these temperatures. Reproducibility and uccuracy. When the recommended procedure was applied to ten replicates containing 1.012 meq of hydrogen peroxide, a standard deviation of 0.005 meq was obtained, with a . . coeffictent of variatton of OS”/,. The relative error of the mean was -0.25%. The calculated Student’s t-value was 2.18, which indicates a significant difference between the experimental mean and the true value. This small error may arise from decomposition of the sample, or the possible mixed potential of the electrode in presence of excess of permanganate might shift the inflexion point of the curve, used as the endpoint, to slightly before the equivalence point.
Determination (NH4
of
the
peroxide
content
in
The proposed procedure was applied to the determination of the peroxide content of the titaniumfluoride-peroxide mixed complex. The contents of Ti, NH;, and F- had been determined previously and the following stoichiometric composition found: Ti:SF:3NHd, which fits the formula (NH4)3Ti02F5. With this stoichiometry, the theoretical content of
I I
the
), Ti02 FS complex
I
1
1
I
75
loo
125
Is0
H202 titrated, % Fig. 3. Effect of temperature on the potentiometric titration curves. HLOr : 1.012 meq; Fe(ll1): 10m4A4; NaF: l.OM; pH = 3.0 KMnO; : 0.0265M.Temperature; (1): 22 ; (2): 30’; (3): 35-.
954
SHORT COMMUNICATIONS
peroxide in the complex would be 13.98”,,. The experimental
value
(mean
13.94”,, with a standard
of
five
determinations)
deviation
4. 1. M. Issa. R. M. Issa. M. A. Khattab and M. R. Mah-
was
of 0.4”,,,. for 0.3-g
samples.
‘. 6.
REFERENCES
1. J. A. Connor and E. A. V. Ebsworth. At/runcrs in Inorgctnic Chrnristr~~c~ntlRadioc~hemistr~, 1964, 6, 280. 2. I. Lange and A. Petzold. Z. Arm/. Chrm.. 1956, 150,24. 3. 1. M. Issa. M. H. Hamdy and A. S. Misbah. Microchrm. J.. 1972. 17. 480.
1. 8. 9. IO.
mond, Intliun J. Chrm., 1975, 13, 942. A. M. Hammam and 1. M. Issa. Mikrochim. kttr. 1976 1, 573. I. M. Issa. M. A. Ghandour and A. M. Hammam. J. IntliunChum. Sot.. 1974. 51. 872. 1. M. Issa, K. A. ldriss and M. M. Ghoneim. Trrltrnf~c. 1976, 23, 249. 1. M. lssa and M. M. Ghoneim. Ttrltrnru, 1973%20, 517. K. A. Idriss, A. M. Hammam. M. M. Seleim and Z. R. El-Komi. Tulunra, 1980, 27, 561. T. J. Lewis. D. H. Richards and D. A. Salter, J. C/wnt. Sot.. 1963. 2434.
r‘,l‘,“r,,. Vol. 2x. pp. 951 I0 954. 1981 Prmted I” Great Br~tam
SHORT
COMMUNICATIONS
POTENTIOMETRIC DETERMINATION OF HYDROGEN PEROXIDE WITH PERMANGANATE IN PRESENCE OF FLUORIDE, WITH IRON AS A CATALYST: APPLICATION TO THE TITANIUM-FLUORIDE-PEROXIDE MIXED COMPLEX L. POLO DIEZ, J. HERNANDEZ MENDEZ and M. J. ALMENDRAL PARRA Department
of Analytical
(Rrcrired
Chemistry.
6 Norrmhrr
Faculty
1980 Reaisrd
of Chemistry.
University
30 April 1981. Accrpred
of Salamanca
Spain
12 May 1981)
Summary-A potentiometric permanganate titration has been developed for the determination of hydrogen peroxide in the presence of fluoride, as well as of the peroxide content in the !itanium-Ruorideperoxide mixed complex. It is based on the stabilization of manganese(II1) with an excess of fluoride in a moderately acidic medium (pH close to 3) and on the use of iron(lIl) as catalyst. Errors are less than 0.5” 0
Several titrimetric methods are available for determination of hydrogen peroxide or the peroxide content of compounds, at macro-levels, those with cerium(IV) or permanganate being the most suitable.’ In the cerium(lV) methods, the presence of fluoride seems to lead to low results,’ but there is no information about its effect on the permanganate methods, although several permangate titrations in fluoride medium have been developed.3m9 In this paper, a potentiometric method is proposed for the determination of hydrogen peroxide with permanganate in the presence of the fluoride, the permanganate being reduced to manganese(M) stabilized as the fluoride complex. The method is applied to the determination of peroxide in the titaniumPfluorideperoxide mixed complex.
the potentials
between
a platinum
wire indicator electrode
and a saturated calomel electrode. Procrdurr Weigh accurately
about
0.1-0.5
g of the titanium-fluori-
de-peroxide complex (or use an equivalent amount of a solution containing this compound or hydrogen peroxide), dissolve it in a small volume of distilled water, and add 10 g of sodium fluoride. Add sufficient 1M sulphuric acid to give a pH close to 3 (Methyl Orange as indicator), enough Fe(lll) solution to give a final concentration about 10e4M. Titrate potentiometrically with the
and of per-
man@nate sohtion. RESULTS AND DISCUSSION Prdiminary
wsults
When the classic permanganate titration in acidic medium was used to determine the peroxide content in the complex, low but reproducible results were obtained. EXPERIMENTAL This error was attributed to partial stabilization of manganese as the manganese(II1) fluoride complex, which was confirmed by the pink colour observed Potassium permanganate solutions. about O.lN. i.e., before the characteristic permanganate colour 0.02M. were standardized against sodium oxalate. Hydroappeared. This fluoride interference cannot be gen peroxide solutions. about O.lN. were standardized avoided by heating in an acidic medium, as this against the permanganate. Other solutions included 1M i brings about decomposition of the peroxide group. and O.lM sulphuric acid. ISM sodium fluoride. and O.OlOM copper sulphate. chromic sulDhate. ammonium To overcome this problem the permanganate titra. metabanadate and ferric sulphate. tion was done in the presence of an excess of fluoride. The best conditions were determined by titrating hyS1o11plr Solid (NH,),Ti02F, was prepared by adding an excess drogen peroxide solutions. The characteristic pink of hydrogen peroxide and ammonium fluoride to a titaof the manganese(IIItfluoride colour complex nium(IV) sulphate solution. which was kept cold and neunecessitated the use of a physical method for the detralized with an ammonia solution to pH 7. termination of the end-point, the potentiometric method with a platinum wire as the indicator elecA Metrohm E-510 potentiometer was used to measure trode being the most suitable.
952
SHORT COMMUNICATIONS
900
-
700-
500-
300 -
100 -
1
t
L
75
loo
I
125
1
150
H*02 titrated, % Fig. I. Efiect of pH on the potentiometric curves. H202: 1.012 meq: NaF: l.OM; Fe(Il1): KMnO,: 0.026SM. pH values are those specifed on the curves.
Porc~llriot?lc~fr;~~
titrcttion
The reaction between permanganate and hydrogen peroxide in the presence of an excess of fluoride at about pH 3 is very slbw. especially near the equivalence point. Moreover. the electrode potential also stabilizes slowly. However. in the presence of copper. vanadium or iron salts. the chemical reaction occurs quickly. as shown by the instantaneous decolorization of permanganate. and the electrode potential stabilizes rapidly. Iron salts produced the best results and were therefore used as the catalyst. According to the literature. iron(III) peroxide species must be involved in the mechanism of the catalysed chemical reaction.1° The potentiometric titration curves obtained in the presence of iron(II1) were well defined, making determination of the end-point easy. A single titration took about IO min. normal for this kind of titration. The potentials obtained with excess of permanganate indicated that the electrode probably responded to a mixed potential. Efl?rr ofpH. Titration curves obtained in the presence of IO- ‘lLf iron(II1) at different pH values are shown in Fig. I. The change in potential around the equivalence point decreases when the pH is increased. and above pH 4 manganese dioxide is formed. The apparent tinal oxidation states of manganese (assum-
ing a quantitative reaction between permanganate and hydrogen peroxide) calculated from the titration curves of Fig. 1 are shown in Table I. These results lead to the conclusion that permanganate is not reduced stoichiometrically to manganese(II) even at pH I. On the other hand, at pH values close to 3. it is reduced quantitatively to manganese(“I). Taking into account that under these conditions the MnF4 complex must predominate in solution, the following reaction presumably occurs MnO;
+ 2H202
+ 4F-
+ 4H+ ---t MnF; + 20,
Titration error assuming the half-reaction MnO; + Mn(II1)
+ 4H,O
As is clear from the table,‘at pH values either lower or higher than 3. other reactions occur simultaneously with the formation of manganese(H) or (IV) respectively. The adjustment of pH to about 3 is easy owing to the presence of the HF/Fbuffer; some Methyl Orange indicator is added, followed by sufficient IM sulphuric acid to turn the colour red. A pH-meter is not needed and a glass electrode should not be used in such a fluoride medium anyway. Ejkt @the iron(lll) concmtration. Concentrations of Fe(II1) in the range 5 x 10-5P10-4M are suitable to make both the chemical and electrochemical reac-
Table I. Apparent tinal oxidation state of manganese and titration error. assuming the MnOJ -+ Mn(IIl) half-reaction at diRerent pH values PH Apparent final oxidation states of Mn
10m4M;
1.0
2.0
3.0
4.0
5.0
2.51
2.96
2.99
3.04
3.77
- 11.0
- I.1
LO.31
+ I.0
+20.3
SHORT
I
I
75
I 150
I 125
I 100
HaOa titrated, %
Fig. 2. EtTect of the NaF concentration on the potentiometric titration curves. Hz02: 0.949 meq: Fe(II1): 10e4M; pH: 3.0: KMn04: 0.0211M. NaF concentration (M); (1) 0.6: (2) 0.8: (3) 1.2 (4) 1.4. tions fast enough. However, as the titration times increased significantly for concentrations lower than 5 x 10-5M, 10e4M was taken as the optimum concentration. Eficr of sodium Juoride concentration. Titration curves obtained at different sodium fluoride concentrations are shown in Fig. 2. The change of potential around the equivalence point decreases slightly for increasing concentrations of sodium fluoride. However. the potential of the MnO;/Mn(lII) system should increase with the sodium fluoride concentration. The cause of the effect observed is obscure, but may be due to the effect of fluoride on the titanium complex as a whole. In other words, the increase in fluoride concentration will increase the concentration of the mixed complex and decrease the
900
953
COMMUNICATIONS
concentration of free peroxide in equilibrium with it, hence increasing the potential of the oxygen/mixedligand complex couple, and accounting for the higher potentials before the end-point. The lower potentials after the end-point could be due to a mixed-potential electrode response. l@ct of temperature. Titration curves obtained at different temperatures are shown in Fig. 3. At room temperatures in the range 15-25” no significant differences are observed in the shape of the titration curves. However, at temperatures higher than about 30”, changes in the titration curves are clear, and low results are obtained. This is probably due to disproportionation of hydrogen peroxide at these temperatures. Reproducibility and uccuracy. When the recommended procedure was applied to ten replicates containing 1.012 meq of hydrogen peroxide, a standard deviation of 0.005 meq was obtained, with a . . coeffictent of variatton of OS”/,. The relative error of the mean was -0.25%. The calculated Student’s t-value was 2.18, which indicates a significant difference between the experimental mean and the true value. This small error may arise from decomposition of the sample, or the possible mixed potential of the electrode in presence of excess of permanganate might shift the inflexion point of the curve, used as the endpoint, to slightly before the equivalence point.
Determination (NH4
of
the
peroxide
content
in
The proposed procedure was applied to the determination of the peroxide content of the titaniumfluoride-peroxide mixed complex. The contents of Ti, NH;, and F- had been determined previously and the following stoichiometric composition found: Ti:SF:3NHd, which fits the formula (NH4)3Ti02F5. With this stoichiometry, the theoretical content of
I I
the
), Ti02 FS complex
I
1
1
I
75
loo
125
Is0
H202 titrated, % Fig. 3. Effect of temperature on the potentiometric titration curves. HLOr : 1.012 meq; Fe(ll1): 10m4A4; NaF: l.OM; pH = 3.0 KMnO; : 0.0265M.Temperature; (1): 22 ; (2): 30’; (3): 35-.
954
SHORT COMMUNICATIONS
peroxide in the complex would be 13.98”,,. The experimental
value
(mean
13.94”,, with a standard
of
five
determinations)
deviation
4. 1. M. Issa. R. M. Issa. M. A. Khattab and M. R. Mah-
was
of 0.4”,,,. for 0.3-g
samples.
‘. 6.
REFERENCES
1. J. A. Connor and E. A. V. Ebsworth. At/runcrs in Inorgctnic Chrnristr~~c~ntlRadioc~hemistr~, 1964, 6, 280. 2. I. Lange and A. Petzold. Z. Arm/. Chrm.. 1956, 150,24. 3. 1. M. Issa. M. H. Hamdy and A. S. Misbah. Microchrm. J.. 1972. 17. 480.
1. 8. 9. IO.
mond, Intliun J. Chrm., 1975, 13, 942. A. M. Hammam and 1. M. Issa. Mikrochim. kttr. 1976 1, 573. I. M. Issa. M. A. Ghandour and A. M. Hammam. J. IntliunChum. Sot.. 1974. 51. 872. 1. M. Issa, K. A. ldriss and M. M. Ghoneim. Trrltrnf~c. 1976, 23, 249. 1. M. lssa and M. M. Ghoneim. Ttrltrnru, 1973%20, 517. K. A. Idriss, A. M. Hammam. M. M. Seleim and Z. R. El-Komi. Tulunra, 1980, 27, 561. T. J. Lewis. D. H. Richards and D. A. Salter, J. C/wnt. Sot.. 1963. 2434.