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Dichromate titration of thallium(I)
Tnlonra, Vol. 22, pp. 93-96. Pergamon
Press
1975. Pmtcd
III Great
Britain
ANNOTATION DICHROMATE TITRATION OF THALLIUM(I) (Received 18 December 1973. Accepted 8 May 1974) The standard redox potential of the thallium(III)-thallium(I) couple is reported’ to be 089 V, which is very nearly the same as that of the chromium(VI)-chromium(III) couple.’ Though this potential is sufficiently high for the absence of oxidation of thallium(I) salts by atmospheric oxygen, several workers have observed the oxidation of thallium(I) salts by atmospheric oxygen in the presence of halide ions. 3-6 Thus the ioss of luminescence of thallium(I) salt solutions under ultraviolet light in the presence of chloride and hydrogen ions was found to be due to the aerial oxidation of the thallium(I) salts4 Oxidation of thallium(I) to thalli~(II1) in aqueous medium at pH less than 6 in the presence of chloride and bromide ions has also been reported5 Daiev et al6 thought it necessary to use an inert atmosphere to prevent aerial oxidation of thallium(I) to thallium(II1) in media containing hydrochloric acid at high concentrations. Rae’ attempted a potentiometric titration of thallium(I) with chromium(V1) but thought that the titration was non-stoichiometric though there is a considerable potential difference between the two redox systems (in media of high hydrochloric acid concentration). According to him the potentiometric titration is not satisfactory even in the presence of iodine mono&Ioride as catalyst. Buds and Erdey’ on the other hand reported a satisfactory potentiometric titration of thallium(I) with chromium(V1) at hydrochloric acid concentration > 5M, but did not find it necessary to use an inert atmosphere for the titration. All the oxidimetri.: titration methods for thallium(I) require the presence of chloride ions9 In view of these findings a study of the reaction between thallium(I) and chromium(W) in hydrochloric acid medium was undertaken.
EXPERIMENTAL
Reagents
The thallium(I) and thallium(II1) solutions were prepared from thallous carbonate and standardized, as described earlier.” Other reagents used were analytical-reagent grade. Apparatus
In the potentiometric titrations the potentials were measured with use of a bright platinum wire indicator electrode and a saturated potassium chloride-agar bridge. During all experiments the thallium solutions were kept under an inert atmosphere, as otherwise there was considerable oxidation of thallium(I) with atmospheric oxygen, especially in solutions containing hydrochloric acid at high concentrations. Formal redox potentials
The formal redox potentials of the thallium(III~thallium(I) couple were measured in solutions of various hydrochloric acid concentration and 0.025M in both thallium(I) and thallium(II1). Air was expelled from these solutions by saturating the solutions with nitrogen and keeping them under a nitrogen atmosphere. The potentials attained by the platinum electrode after 10-15 mitt were stable and were recorded. A mixture of 50 ml of concentrated hydrochloric acid and 40 ml of water in a 150-ml titration vessel is saturated with carbon dioxide by passage of the gas for a few minutes. An aliquot of thallium(I) solution is added and the potential acquired by a bright platinum electrode inserted in the mixture is measured against a saturated calomel electrode. The potential is then recorded after each addition of titrant and the equivalence-point is found in the usual way. Reasonably stable potentials are attained rapidly by the indicator electrode until near the equivalence point. At the equivalence point the stabilization of potential is rather slow, a few minutes wait being necessary before the measurement. The potential break at the equivalence point is about 150 mV for the addition of @05 ml of 0.05N chromium(V1). Among other methods of detecting the end-point of this titration the extract& end-point method of Rae” and the chemiluminescent method of BtuSs and Erdey’ may be mentioned. No reversible redox indicator seems to have been reported. Ferroin, which is widely used in titrations with cerium(IVk has a formal oxidation potential of 1.06 V,” and might function as a reversible indicator in the present titration. Experiments have shown
ANNOTATION
94
Table 1. Determination of thallium(I) by titration with potassium dichromate in 6A4 hydrochloric acid medium Method of detecting the end-point
Thallium(I), mmole Taken Found 00844 0.1688 0.2532 03376 0.4220 0.6330
@0848 0.1682 0.2524 0.3372 0.4208 0.6300
O-0485 0.1309 0.1455 0.1746 0.1940 0.2425
OG483 @1315 01448 0.1744 0.1932 0.2423
Potentiometry
that ferroin is rapidly oxidized by dichromate, and ferriin is rapidly reduced by thallium(I) in media > 5M in hydrochlqric acid, and the indictor gives satisfactory results (Table 1). Tartrate and iron interfere in this titration of thallium(I) with dichromate. Other reducing agents which react with dichromate r&t also be absent. RESULTS
AND
DISCUSSION
The formal redox potentials of the thallium(III)-thallium(I) couple, as determined in the present study, and those reported’ for the chromium(VI)-chromium(III) couple, in media of various hydrochloric acid concentration, are presented in Fig. 1. From these results it is clear that while the formal potentials of the chromium system increase, those of the thallium system decrease with increase in the concentration of hydrochloric acid. Thus though the formal redox potentials of the two systems are comparable at low acidity, the difference between them increases considerably at high hydrochloric acid concentration, allowing a satisfactory oxidimetric titration of thallium(I) with chromium(W). Besides the thermodynamics, the kinetics of the reaction must also be favourable if the titration is to be satisfactory. The rates of the reduction of chromium(W) by thallium(I) were studied spectrophotometrically. The absorption spectra of chromium(V1) and chromium(II1) between 340 and 600 nm do not vary much with variation in
-
CR (XI) -CR (ml system ~TtmI,-TlcI, system
II I II I 460 0 I2 3 4 5 6 Conc%ntratlond hydrochloric acid,
II 7 6 molorlty
I 9
/ IO
Fig. 1. Formal redox potentials of TI(III)/Tl(I) and Cr(VI)/Cr(III) systems in hydrochloric acid media.
95
ANNOTATION
0.3 -05MHCl e I.OM WI -2.OM HCL -3OMHCI -4.OM HCI -5OM HCl
0.2
o
to
20
3040
so 60 708090 Tlfm. min
Ii0 120
Fig. 2. Effect of hydrochloric acid concentration on the kinetics of the reaction between Tl(1) and Cr(V1). hydrochloric acid concentration from 1 to 6M. Hence the reaction rate at room temperature (29”) was determined by measuring the absorbance (at 360 nm) of mixtures in various acid media, as shown in Fig. 2. The results show that (if the speed of the reaction between thallium(I) and chromium(V1) increases with the concentration of the hydrochloric acid, (iQ the reaction does not appear to take place if the medium is MM or less in hydrochloric acid, (iii) the reaction is quite rapid, complete and comes to equilibrium in less than 2 mm if the hydrochloric acid concentrationis > 5M. Similar experiments were also carried out at go*, the volume of the mixture being kept constant and the mixture not in contact with a~osph~ic oxygen. The reaction is then complete within 3 min, provided the hydrochloric acid concentration is 3 3&f. When the acid is only 2&f, the reaction is not complete even after, heating for more than an hour. From these observations it is clear that the titration of thallium(I) with potassium dicbromate should be feasible provided the medium is > 5M in hydrochloric acid. The results obtained by application of the procedure are given in Table 1. That the effect of hydrochloric acid concentration on the potential of the thallium couple is due to the chloride ion was established by keeping the hydrogen ion concentration constant with percbloric acid and varying the chloride concentration, and by varying the acidity in the absence of chloride (Tables 2 and 3). However, according to the thermodynamics the titration should be feasible in 2M hydrochloric acid, and the necessity to use a much Table 2. Formal redox potentials of Tl(III)/n(I) system in mineral acids (us. N.H.E.), mV Concentration of acid, M 0.10 025 050 1.00 250 5.00 750 1000 11.85 12.50 15GO
HClOz, 1236 1255 1249 1249 1272 1330 1433 -
HN03
HISO*
1279 1267 1252 1240 1224 1221 1230 1248 1270 -
1215 1212 1210 1216 1225 1230 1273 1320 1372 1427
*H,PO* -I-0.5M H$O, 1205 1197 1188 1179 1166 1167 1178 1195 1204 -
HCl 862 816 782 746 682 606 553 513 -
* Since t~lliu~II1) hydroxide does not dissolve easily in phosphoric acid it is dissolved in minimum amount of sulphuric acid before addition to phosphoric acid of the required strength.
ANNOTATION
96
Table 3. Effect of varying con~ntrations of hydrogen ion and chloride ion on the formal redox potentials of the ~(III)~l(I) system (us. NILE.), mV
Ha% M
1.00 1.00 1.00 1.00 1.00 1.00 1GO .5+K-l 5+0 5.00 500 500 500
Chloride, M
&025 0.01 O*lO 0.50 1.00 2.00 0 0.0025 @Ol 0.10 050 lG0
Formal redox potential of Tl(III)/Tl(I)
Chloride,
1250 1156 989 867 784 768 721 1330 1266 967 820 743 721
M
HaO.+, M
Formal redox potential of Tl(III)/Tl(I)
@1 0.1 0.1 O-1 O-1 0.1 0.1 1.0 1.0 1.0 1.0 1.0 1.0 1-O
0.10 0.50 1.00 2.00 3.00 4.00 500 0.10 O-SO I.00 z-00 3.00 4.00 500
883 882 867 855 844 831 821 785 774 768 7.53 738 726 721
0.10 0.25 075 1M) 1.25 1.75
1.90 1.75 1.25 1.00 0.75 0.25
854 844 794 768 766 762
higher acidity implies that hydrogen ions play an important part in determiuing the kinetics (as would be expected from the equation for the reaction). As there is no excess of dichromate in the titration solution until after the end-point there will be no oxidation of chloride by the dichromate.
Department ofChemistry Andhra University Waltair, India
S. R. SAGI C. S. PRAKASARAJU K.V. RAMANA REFERENCES
1. R. H. Hughes and C. S. Garner, J. Am. Chem. Sot., 1942,64,1644. 2. G. F. Smith and F. P. Richter, Ind. Eng. Chem., Anal. Ed., 1944,16,580. 3. G. F. Kirkbright, P. J. Mayne and T. S. West, J. Chem. Sot., Dalton Trans., 1972,17, 1918. 4. M. V. Billi and B. A. Okhrinenko, Visn. Kiius’k. Univ., Ser. Astron., Fiz. ta Khim., 1962, No. 5, 15. 5. Teijin Ltd., N&h, Pat. A&., 6,505,487 (cl Co 1 g) 2 Nov. 1965, Japan Pat. Appl. 1 May 1964 and 19 Jan. 1965. 6. Kh. Daiev and S. Nikita, Godjsh~ik So~skiya Univ. Fir. Mat. Fak. Kniga 3 -Khim., 196566,6& 33. 7. K. B. Rao, I)..%. Thesis, Andhra Univ., Waltair, India, 1957. 8. I. Buz&s and L. Erdey, Talanta, 1963, 10,467. 9. I. M. Kolthoff and R. Belcher, Volumetric Analysis, Vol. HI, pp. 102, 103, 145, 453,461, 519, 641. Interscience, New York, 19.57. 10. S. R. Sagi and K. V. Ramana, Talanta, 1969,16,1217. 11. K. B. Rao, 2. A&. Chem., 1959,165, 193. 12. I. M. Kolthoff and D. N. Hume, J. Am. Chem. SOL, 1943.65, 1895.
Summary-The formal redox potentials of the thallium(III+allium(I) couple in different acids of varying strengths are reported. The minimum concentration of hydrochloric acid required for a direct titration of thallium(I) with potassium dichromate is SM. Thallium(I) can be titrated directly with the primary standard oxidant, potassium dichromate, at room temperature, with ferroin as indicator, in 6M hydrocNoric acid. Atmospheric oxygen must be excluded.
Press
1975. Pmtcd
III Great
Britain
ANNOTATION DICHROMATE TITRATION OF THALLIUM(I) (Received 18 December 1973. Accepted 8 May 1974) The standard redox potential of the thallium(III)-thallium(I) couple is reported’ to be 089 V, which is very nearly the same as that of the chromium(VI)-chromium(III) couple.’ Though this potential is sufficiently high for the absence of oxidation of thallium(I) salts by atmospheric oxygen, several workers have observed the oxidation of thallium(I) salts by atmospheric oxygen in the presence of halide ions. 3-6 Thus the ioss of luminescence of thallium(I) salt solutions under ultraviolet light in the presence of chloride and hydrogen ions was found to be due to the aerial oxidation of the thallium(I) salts4 Oxidation of thallium(I) to thalli~(II1) in aqueous medium at pH less than 6 in the presence of chloride and bromide ions has also been reported5 Daiev et al6 thought it necessary to use an inert atmosphere to prevent aerial oxidation of thallium(I) to thallium(II1) in media containing hydrochloric acid at high concentrations. Rae’ attempted a potentiometric titration of thallium(I) with chromium(V1) but thought that the titration was non-stoichiometric though there is a considerable potential difference between the two redox systems (in media of high hydrochloric acid concentration). According to him the potentiometric titration is not satisfactory even in the presence of iodine mono&Ioride as catalyst. Buds and Erdey’ on the other hand reported a satisfactory potentiometric titration of thallium(I) with chromium(V1) at hydrochloric acid concentration > 5M, but did not find it necessary to use an inert atmosphere for the titration. All the oxidimetri.: titration methods for thallium(I) require the presence of chloride ions9 In view of these findings a study of the reaction between thallium(I) and chromium(W) in hydrochloric acid medium was undertaken.
EXPERIMENTAL
Reagents
The thallium(I) and thallium(II1) solutions were prepared from thallous carbonate and standardized, as described earlier.” Other reagents used were analytical-reagent grade. Apparatus
In the potentiometric titrations the potentials were measured with use of a bright platinum wire indicator electrode and a saturated potassium chloride-agar bridge. During all experiments the thallium solutions were kept under an inert atmosphere, as otherwise there was considerable oxidation of thallium(I) with atmospheric oxygen, especially in solutions containing hydrochloric acid at high concentrations. Formal redox potentials
The formal redox potentials of the thallium(III~thallium(I) couple were measured in solutions of various hydrochloric acid concentration and 0.025M in both thallium(I) and thallium(II1). Air was expelled from these solutions by saturating the solutions with nitrogen and keeping them under a nitrogen atmosphere. The potentials attained by the platinum electrode after 10-15 mitt were stable and were recorded. A mixture of 50 ml of concentrated hydrochloric acid and 40 ml of water in a 150-ml titration vessel is saturated with carbon dioxide by passage of the gas for a few minutes. An aliquot of thallium(I) solution is added and the potential acquired by a bright platinum electrode inserted in the mixture is measured against a saturated calomel electrode. The potential is then recorded after each addition of titrant and the equivalence-point is found in the usual way. Reasonably stable potentials are attained rapidly by the indicator electrode until near the equivalence point. At the equivalence point the stabilization of potential is rather slow, a few minutes wait being necessary before the measurement. The potential break at the equivalence point is about 150 mV for the addition of @05 ml of 0.05N chromium(V1). Among other methods of detecting the end-point of this titration the extract& end-point method of Rae” and the chemiluminescent method of BtuSs and Erdey’ may be mentioned. No reversible redox indicator seems to have been reported. Ferroin, which is widely used in titrations with cerium(IVk has a formal oxidation potential of 1.06 V,” and might function as a reversible indicator in the present titration. Experiments have shown
ANNOTATION
94
Table 1. Determination of thallium(I) by titration with potassium dichromate in 6A4 hydrochloric acid medium Method of detecting the end-point
Thallium(I), mmole Taken Found 00844 0.1688 0.2532 03376 0.4220 0.6330
@0848 0.1682 0.2524 0.3372 0.4208 0.6300
O-0485 0.1309 0.1455 0.1746 0.1940 0.2425
OG483 @1315 01448 0.1744 0.1932 0.2423
Potentiometry
that ferroin is rapidly oxidized by dichromate, and ferriin is rapidly reduced by thallium(I) in media > 5M in hydrochlqric acid, and the indictor gives satisfactory results (Table 1). Tartrate and iron interfere in this titration of thallium(I) with dichromate. Other reducing agents which react with dichromate r&t also be absent. RESULTS
AND
DISCUSSION
The formal redox potentials of the thallium(III)-thallium(I) couple, as determined in the present study, and those reported’ for the chromium(VI)-chromium(III) couple, in media of various hydrochloric acid concentration, are presented in Fig. 1. From these results it is clear that while the formal potentials of the chromium system increase, those of the thallium system decrease with increase in the concentration of hydrochloric acid. Thus though the formal redox potentials of the two systems are comparable at low acidity, the difference between them increases considerably at high hydrochloric acid concentration, allowing a satisfactory oxidimetric titration of thallium(I) with chromium(W). Besides the thermodynamics, the kinetics of the reaction must also be favourable if the titration is to be satisfactory. The rates of the reduction of chromium(W) by thallium(I) were studied spectrophotometrically. The absorption spectra of chromium(V1) and chromium(II1) between 340 and 600 nm do not vary much with variation in
-
CR (XI) -CR (ml system ~TtmI,-TlcI, system
II I II I 460 0 I2 3 4 5 6 Conc%ntratlond hydrochloric acid,
II 7 6 molorlty
I 9
/ IO
Fig. 1. Formal redox potentials of TI(III)/Tl(I) and Cr(VI)/Cr(III) systems in hydrochloric acid media.
95
ANNOTATION
0.3 -05MHCl e I.OM WI -2.OM HCL -3OMHCI -4.OM HCI -5OM HCl
0.2
o
to
20
3040
so 60 708090 Tlfm. min
Ii0 120
Fig. 2. Effect of hydrochloric acid concentration on the kinetics of the reaction between Tl(1) and Cr(V1). hydrochloric acid concentration from 1 to 6M. Hence the reaction rate at room temperature (29”) was determined by measuring the absorbance (at 360 nm) of mixtures in various acid media, as shown in Fig. 2. The results show that (if the speed of the reaction between thallium(I) and chromium(V1) increases with the concentration of the hydrochloric acid, (iQ the reaction does not appear to take place if the medium is MM or less in hydrochloric acid, (iii) the reaction is quite rapid, complete and comes to equilibrium in less than 2 mm if the hydrochloric acid concentrationis > 5M. Similar experiments were also carried out at go*, the volume of the mixture being kept constant and the mixture not in contact with a~osph~ic oxygen. The reaction is then complete within 3 min, provided the hydrochloric acid concentration is 3 3&f. When the acid is only 2&f, the reaction is not complete even after, heating for more than an hour. From these observations it is clear that the titration of thallium(I) with potassium dicbromate should be feasible provided the medium is > 5M in hydrochloric acid. The results obtained by application of the procedure are given in Table 1. That the effect of hydrochloric acid concentration on the potential of the thallium couple is due to the chloride ion was established by keeping the hydrogen ion concentration constant with percbloric acid and varying the chloride concentration, and by varying the acidity in the absence of chloride (Tables 2 and 3). However, according to the thermodynamics the titration should be feasible in 2M hydrochloric acid, and the necessity to use a much Table 2. Formal redox potentials of Tl(III)/n(I) system in mineral acids (us. N.H.E.), mV Concentration of acid, M 0.10 025 050 1.00 250 5.00 750 1000 11.85 12.50 15GO
HClOz, 1236 1255 1249 1249 1272 1330 1433 -
HN03
HISO*
1279 1267 1252 1240 1224 1221 1230 1248 1270 -
1215 1212 1210 1216 1225 1230 1273 1320 1372 1427
*H,PO* -I-0.5M H$O, 1205 1197 1188 1179 1166 1167 1178 1195 1204 -
HCl 862 816 782 746 682 606 553 513 -
* Since t~lliu~II1) hydroxide does not dissolve easily in phosphoric acid it is dissolved in minimum amount of sulphuric acid before addition to phosphoric acid of the required strength.
ANNOTATION
96
Table 3. Effect of varying con~ntrations of hydrogen ion and chloride ion on the formal redox potentials of the ~(III)~l(I) system (us. NILE.), mV
Ha% M
1.00 1.00 1.00 1.00 1.00 1.00 1GO .5+K-l 5+0 5.00 500 500 500
Chloride, M
&025 0.01 O*lO 0.50 1.00 2.00 0 0.0025 @Ol 0.10 050 lG0
Formal redox potential of Tl(III)/Tl(I)
Chloride,
1250 1156 989 867 784 768 721 1330 1266 967 820 743 721
M
HaO.+, M
Formal redox potential of Tl(III)/Tl(I)
@1 0.1 0.1 O-1 O-1 0.1 0.1 1.0 1.0 1.0 1.0 1.0 1.0 1-O
0.10 0.50 1.00 2.00 3.00 4.00 500 0.10 O-SO I.00 z-00 3.00 4.00 500
883 882 867 855 844 831 821 785 774 768 7.53 738 726 721
0.10 0.25 075 1M) 1.25 1.75
1.90 1.75 1.25 1.00 0.75 0.25
854 844 794 768 766 762
higher acidity implies that hydrogen ions play an important part in determiuing the kinetics (as would be expected from the equation for the reaction). As there is no excess of dichromate in the titration solution until after the end-point there will be no oxidation of chloride by the dichromate.
Department ofChemistry Andhra University Waltair, India
S. R. SAGI C. S. PRAKASARAJU K.V. RAMANA REFERENCES
1. R. H. Hughes and C. S. Garner, J. Am. Chem. Sot., 1942,64,1644. 2. G. F. Smith and F. P. Richter, Ind. Eng. Chem., Anal. Ed., 1944,16,580. 3. G. F. Kirkbright, P. J. Mayne and T. S. West, J. Chem. Sot., Dalton Trans., 1972,17, 1918. 4. M. V. Billi and B. A. Okhrinenko, Visn. Kiius’k. Univ., Ser. Astron., Fiz. ta Khim., 1962, No. 5, 15. 5. Teijin Ltd., N&h, Pat. A&., 6,505,487 (cl Co 1 g) 2 Nov. 1965, Japan Pat. Appl. 1 May 1964 and 19 Jan. 1965. 6. Kh. Daiev and S. Nikita, Godjsh~ik So~skiya Univ. Fir. Mat. Fak. Kniga 3 -Khim., 196566,6& 33. 7. K. B. Rao, I)..%. Thesis, Andhra Univ., Waltair, India, 1957. 8. I. Buz&s and L. Erdey, Talanta, 1963, 10,467. 9. I. M. Kolthoff and R. Belcher, Volumetric Analysis, Vol. HI, pp. 102, 103, 145, 453,461, 519, 641. Interscience, New York, 19.57. 10. S. R. Sagi and K. V. Ramana, Talanta, 1969,16,1217. 11. K. B. Rao, 2. A&. Chem., 1959,165, 193. 12. I. M. Kolthoff and D. N. Hume, J. Am. Chem. SOL, 1943.65, 1895.
Summary-The formal redox potentials of the thallium(III+allium(I) couple in different acids of varying strengths are reported. The minimum concentration of hydrochloric acid required for a direct titration of thallium(I) with potassium dichromate is SM. Thallium(I) can be titrated directly with the primary standard oxidant, potassium dichromate, at room temperature, with ferroin as indicator, in 6M hydrocNoric acid. Atmospheric oxygen must be excluded.