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Interhaloid complexes in aqueous solution
J. Inorg. Nucl. Chem., 1959, Vol. 11, pp. 56 to 61. Pergamon Press Ltd. Printed in Northern Ireland
INTERHALOID COMPLEXES IN AQUEOUS SOLUTION E. PUNGOR, K. BrinGER and E. SCrr~LEK Institute of Inorganic and Analytical Chemistry, L. E6tv6s University, Budapest (Received 14 July 1958; in revised.form 14 October 1958)
Abstract--The examination of a solution of bromine chloride in hydrochloric acid by ultra-violet absorption and redox potential measurements has demonstrated the existence in these solutions of a complex ion of formula BrCI6~-, with a stability constant of (2.6 ~: 1"0) × 10~. The formulae of the similar complex ions formed from iodine chloride and bromide were shown to be ICI6~- and IBr43with stability constants (4-3 4- 2-2) × 102 and '(5"0 -4- 1"0) × 10s. The experimental procedure was shown to be correct by an investigation of the Iv- complex. IT is knowntX, 2~ that, of the interhaloid compounds, iodine monochloride and iodine monobromide can form interhaloid complex anions in aqueous solution. ICl 2and IBr 2- have been suggested as the formulae of these complexes, the central unipositive iodine having two chloride or bromide ions as ligands, chlorine and bromine having higher electron affinities than iodine. By studying partition equilibria in water-chloroform, FAULL~1) has found dissociation constants of 6 × 10-a and 2.7 × 10-3 respectively for IC12- and IBr2-. N o reference has been found in the literature to the formation of a similar complex anion from BrC1. Solutions of bromine chloride, when stored, gradually become more dilute because of the halogen volatility. But during an investigation of bromination of aromatic compounds by bromine chloride taJ it was found that the fall in concentration of a 0.1 N solution did not exceed 3-5 per cent in 2-3 months. It was suggested that the comparative stability of such solutions was due to the formation of a complex chloride of bromine chloride. In the present work, we have investigated the constitution and stability constant of this presumed complex ion, and of the chloride complex o f iodine chloride, and the bromide complex of iodine bromide, discussed above. Ultra-violet absorption measurements were made using a Beckman D.U. quartz spectrophotometer, with a hydrogen lamp and 1 cm quartz cells. Redox potentials were measured with an electronic p H meter Radiometer 22, using 20 × 25 mm smooth platinum and saturated calomel electrodes. RESULTS 1. The ultra-violet absorption spectra of "neutral" solutions of bromine chloride containing no acid are characterized by extremely flat maxima. When sulphuric acid is added to the solutions, a more distinct, though still flat, maximum appears. However, in the presence of chloride ions, a characteristic sharp maximum is obtained at 343 m/~ (Fig.l). The extinction coefficient at maximum absorption in the presence of sulphuric acid Ill j. H. FAULL,Jr. Amer. Chem. Soc. 56, 522 (1934). ~2~G. S. FogaES,S. W. GLASSand R. M. Fuoss, J. Amer. Chem. Soc. 47, 2892 (1925). ~al E. SCHULEKand K. BURGER,Talanta 1, 219 (1958). 56
lnterhaloid complexes in aqueous solution
57
was found to be independent of the concentration of sulphuric acid (Fig. 2), but in chloride solutions, it increases with the concentration of chloride ion (Fig. 3). 2. The redox potentials of bromine chloride solutions decreased as the concentration of chloride ion was increased. Increase in the halide concentration also decreased the redox potentials of solutions of iodine chloride in hydrochloric acid and iodine bromide in bromide ion. 19E 1,z 4
4
4"t 4"0
2 0'3 300
350
400 m)~,
Fl(;. l.--Ultra-violet absorption curves of solutions of bromine chloride. 1. 0.005 M solution of BrCI in 1.5 N HC1 2. 0.005 M solution of BrCI in 0.5 N HCI 0.005 M solution of BrCI in 0.5 N NaCI 3. 0.005 M solution of BrCI in 0'2 N NaCI 4. 0.005 M solution of BrCI in 1.0 N H2SO4 5. 0-005 M solution of BrCI in water.
~'5
I"# / '1.3
E (343 m~)
'1.'2
'
4,4
0,7
|
!
I
!
I
1"0 ~.'0 3'0 ~'0 5"0 N N~,i, FIG. 2.--Effect of concentration of sulpburic acid on the extinction values of bromine chloride. E343p in 1.0, 2.0, 3.0, 4-0, 5.0 N HzSO4.
FIG. L---Effect of concentration of hydrochloric acid on the extinction values of bromine chloride. Es43t~ in 1.0, 2.0, 3.0, 4.0, 5.0, N HCI.
58
E. PUNGOR, K. BURGER and E. SCHULEK
THEORETICAL INTERPRETATION The above results confirm the formation of a complex BrCI!~x-l~j- from BrCl and chloride ion. From the variation of the redox potential of iodine chloride and bromide solutions with the corresponding halide ion concentration, we have attempted to calculate the stability constants and co-ordination numbers of these complexes, the published values of which were calculated from partition equilibria in water-chloroform, where hydration may interfere to a considerable extent. The redox potential of an interhaloid XY is governed by the concentration [XY] of the actual interhaloid present, according to the equation 0.058 E -- E o + - - T - l o g ([XY]/[X-] [Y-]) (1) and we have assumed in our calculation that the anion complex XYIf"- 1)1 affects the redox potential only indirectly by changing [XY]. Owing to the limited number of measurements possible, we have not been able to elucidate individual stability constants, but only the overall constant. That, is, the K we have calculated is the product of the individual constants for the consecutive dissociations. When the necessary substitutions have been made, we find ( E--
Eo
0.058 2 log [X-]
)O.O,8 = -- 2
log {[XYtot]/([Y l +
K[KY-]")}
(2)
where brackets represent activities (the activity coefficients being calculated from the Debye-Htickel equation), n is the co-ordination number and [XYtot] the total XY concentration. Since in the present work, the stoicheiometric interhaloid, containing no halide 0.058 ions of the more positive element, were used, the term E 0 - T log [X-] in equation (2) remains constant, and may be denoted by K. We may therefore substitute the measured e.m.f.'s into 0.058 E -- X = T log {[XYtot]/([Y-] + K[Y-]")} (3) and solve these equations for n, the co-ordination number, and Kthe stability constant of the complex in question. These values were then used to give calculated e.m.f. values to compare with the measured values--see Tables I-3. The good agreement between these values confirms the assumption that the redox potentials are determined by the concentration of free interhaloid present at equilibrium. Our results indicate the following formulae, co-ordination numbers and stability constants for the complexes investigated. Complex formula
6
n
K-
BrC1 in CI] IC1 in ClIBr in Br]BrC165-,or BrCI (HC1)51 1C16~-, or ICI(HCI) 5 IBr4a-, or IBr (HBr)z
-
Published value for K
1.0 x
6
4
t
(4.3 zk 2"2) × 102
(5"0 ± 1.0) × 103
L
1.7 × 102
3.7 × 102
Interhaloid complexes in aqueous solution
59
TABLE I.---EFFECT OF CONCENTRATION OF HYDROCHLORIC ACID ON rile REDOX POFENT1AL OF A SOLUTIONOF BROMINECHLORIDE IN HYDROCHLORIC ACID Redox potentials BrCI (mole)
HCI (mole)
0"005
0"05
TABLE 2 . - - E F F E C T
measured (mY)
calculated (mY)
0"2 0"5 I'0 1.5 2'0 2'5 3"0 4"0 5"0
1323 1290 1268 1237 1230 1209 1195 1170 1156
1317 1300 1263 1239 1219 1206 1191 1175 1154
0.6 1'0 1'5 2"0 2"5 3'0
1356 1336 1309 1283 1267 1253
1361 1332 1308 1288 1276 1260
OF CONCENTRATION OF HYDROCHLORIC ACID ON THE REDOX
POTENTIAL OF A SOLUTION OF IODINE CHLORIDE IN HYDROCHLORIC ACID
Redox potentials ICI (mole)
0.01
0'02
HCI (mole)
0'2 0"5 1"0 I '5 2'0 2"5 3"0
0"2 0"5 I'0 1"5 2"0 2"5 3"0
i
mesaured (mY)
calculated (mV)
1060 1040 1000 979 962 948 935
1064 1044 1005 978 957 944 931
1080 1050 1018 997 984 970 956
1084 1063 1024 997 977 963 953
60
E. PuNc_,oR, K. BUROEtXand E. Scm.~..Eg:
TABLE 3.--E~ECT
o F CONCENTRATION OF BROMIDE IONS ON ~
REDOX POTENTIAL
OF A SOLUTION OF IODINE BROMIDE CONTAINING BROMIDE IONS I
Rcdox potentials IBr (mole)
0"01
• 0"02
Br(mole)
measured (mV)
calculated (mV)
0-5 1.0 1"5 2"0 2'5 3"0
860 830 815 805 792 781
857 827 810 796 788 780
0"2 0"5 1-0 1"5 2.0 2-5 3"0
925 880 850 832 823 809 801
921 881 852 834 820 811 804
TABLE 4 . - - E F F E C T OF CONCENTRATION OF IODIDE IONS ON THE REDOX POTENTIAL OF A SOLUTION OF IODINE CONTAINING IODIDE IONS
Rcdox potentials Is
I-
(mole)
(mole)
measured (mY)
calculated (mV)
0"01
0"5 l'O 1"5 2"0 2"5
490 468 459 447 44O
479 465 456 449 444
0.02
0.5 1.0 1"5 2"0 2-5
503 481 469 460 456
497 482 473 466 462
lnterhaloitl complexes in aqueous solution
61
Since our calculations for both iodine chloride and bromide gave values of K and n differing from those given in the literature, it seemed necessary to verify the validity of our experimental method. Investigation of the tri-iodide complex seemed suitable for this purpose, the co-ordination number and stability constant having been calculated in a perfectly reliable manner. ~41 From redox potential measurements, followed by the use of equation (3), these parameters were calculated for the tri-iodide complex; the co-ordination number was found to be two, and the stability constant (3.2-q- 1.3) ~'< 102, (Table 4) in fair agreement with the published figure of 7 × 102. The correctness of the assumption involved in our method is thus demonstrated. '~ G. Jo~Es a n d B. B.
KAPLAN,J. Amer. Chem. Soc. 50, 1845 (1928).
INTERHALOID COMPLEXES IN AQUEOUS SOLUTION E. PUNGOR, K. BrinGER and E. SCrr~LEK Institute of Inorganic and Analytical Chemistry, L. E6tv6s University, Budapest (Received 14 July 1958; in revised.form 14 October 1958)
Abstract--The examination of a solution of bromine chloride in hydrochloric acid by ultra-violet absorption and redox potential measurements has demonstrated the existence in these solutions of a complex ion of formula BrCI6~-, with a stability constant of (2.6 ~: 1"0) × 10~. The formulae of the similar complex ions formed from iodine chloride and bromide were shown to be ICI6~- and IBr43with stability constants (4-3 4- 2-2) × 102 and '(5"0 -4- 1"0) × 10s. The experimental procedure was shown to be correct by an investigation of the Iv- complex. IT is knowntX, 2~ that, of the interhaloid compounds, iodine monochloride and iodine monobromide can form interhaloid complex anions in aqueous solution. ICl 2and IBr 2- have been suggested as the formulae of these complexes, the central unipositive iodine having two chloride or bromide ions as ligands, chlorine and bromine having higher electron affinities than iodine. By studying partition equilibria in water-chloroform, FAULL~1) has found dissociation constants of 6 × 10-a and 2.7 × 10-3 respectively for IC12- and IBr2-. N o reference has been found in the literature to the formation of a similar complex anion from BrC1. Solutions of bromine chloride, when stored, gradually become more dilute because of the halogen volatility. But during an investigation of bromination of aromatic compounds by bromine chloride taJ it was found that the fall in concentration of a 0.1 N solution did not exceed 3-5 per cent in 2-3 months. It was suggested that the comparative stability of such solutions was due to the formation of a complex chloride of bromine chloride. In the present work, we have investigated the constitution and stability constant of this presumed complex ion, and of the chloride complex o f iodine chloride, and the bromide complex of iodine bromide, discussed above. Ultra-violet absorption measurements were made using a Beckman D.U. quartz spectrophotometer, with a hydrogen lamp and 1 cm quartz cells. Redox potentials were measured with an electronic p H meter Radiometer 22, using 20 × 25 mm smooth platinum and saturated calomel electrodes. RESULTS 1. The ultra-violet absorption spectra of "neutral" solutions of bromine chloride containing no acid are characterized by extremely flat maxima. When sulphuric acid is added to the solutions, a more distinct, though still flat, maximum appears. However, in the presence of chloride ions, a characteristic sharp maximum is obtained at 343 m/~ (Fig.l). The extinction coefficient at maximum absorption in the presence of sulphuric acid Ill j. H. FAULL,Jr. Amer. Chem. Soc. 56, 522 (1934). ~2~G. S. FogaES,S. W. GLASSand R. M. Fuoss, J. Amer. Chem. Soc. 47, 2892 (1925). ~al E. SCHULEKand K. BURGER,Talanta 1, 219 (1958). 56
lnterhaloid complexes in aqueous solution
57
was found to be independent of the concentration of sulphuric acid (Fig. 2), but in chloride solutions, it increases with the concentration of chloride ion (Fig. 3). 2. The redox potentials of bromine chloride solutions decreased as the concentration of chloride ion was increased. Increase in the halide concentration also decreased the redox potentials of solutions of iodine chloride in hydrochloric acid and iodine bromide in bromide ion. 19E 1,z 4
4
4"t 4"0
2 0'3 300
350
400 m)~,
Fl(;. l.--Ultra-violet absorption curves of solutions of bromine chloride. 1. 0.005 M solution of BrCI in 1.5 N HC1 2. 0.005 M solution of BrCI in 0.5 N HCI 0.005 M solution of BrCI in 0.5 N NaCI 3. 0.005 M solution of BrCI in 0'2 N NaCI 4. 0.005 M solution of BrCI in 1.0 N H2SO4 5. 0-005 M solution of BrCI in water.
~'5
I"# / '1.3
E (343 m~)
'1.'2
'
4,4
0,7
|
!
I
!
I
1"0 ~.'0 3'0 ~'0 5"0 N N~,i, FIG. 2.--Effect of concentration of sulpburic acid on the extinction values of bromine chloride. E343p in 1.0, 2.0, 3.0, 4-0, 5.0 N HzSO4.
FIG. L---Effect of concentration of hydrochloric acid on the extinction values of bromine chloride. Es43t~ in 1.0, 2.0, 3.0, 4.0, 5.0, N HCI.
58
E. PUNGOR, K. BURGER and E. SCHULEK
THEORETICAL INTERPRETATION The above results confirm the formation of a complex BrCI!~x-l~j- from BrCl and chloride ion. From the variation of the redox potential of iodine chloride and bromide solutions with the corresponding halide ion concentration, we have attempted to calculate the stability constants and co-ordination numbers of these complexes, the published values of which were calculated from partition equilibria in water-chloroform, where hydration may interfere to a considerable extent. The redox potential of an interhaloid XY is governed by the concentration [XY] of the actual interhaloid present, according to the equation 0.058 E -- E o + - - T - l o g ([XY]/[X-] [Y-]) (1) and we have assumed in our calculation that the anion complex XYIf"- 1)1 affects the redox potential only indirectly by changing [XY]. Owing to the limited number of measurements possible, we have not been able to elucidate individual stability constants, but only the overall constant. That, is, the K we have calculated is the product of the individual constants for the consecutive dissociations. When the necessary substitutions have been made, we find ( E--
Eo
0.058 2 log [X-]
)O.O,8 = -- 2
log {[XYtot]/([Y l +
K[KY-]")}
(2)
where brackets represent activities (the activity coefficients being calculated from the Debye-Htickel equation), n is the co-ordination number and [XYtot] the total XY concentration. Since in the present work, the stoicheiometric interhaloid, containing no halide 0.058 ions of the more positive element, were used, the term E 0 - T log [X-] in equation (2) remains constant, and may be denoted by K. We may therefore substitute the measured e.m.f.'s into 0.058 E -- X = T log {[XYtot]/([Y-] + K[Y-]")} (3) and solve these equations for n, the co-ordination number, and Kthe stability constant of the complex in question. These values were then used to give calculated e.m.f. values to compare with the measured values--see Tables I-3. The good agreement between these values confirms the assumption that the redox potentials are determined by the concentration of free interhaloid present at equilibrium. Our results indicate the following formulae, co-ordination numbers and stability constants for the complexes investigated. Complex formula
6
n
K-
BrC1 in CI] IC1 in ClIBr in Br]BrC165-,or BrCI (HC1)51 1C16~-, or ICI(HCI) 5 IBr4a-, or IBr (HBr)z
-
Published value for K
1.0 x
6
4
t
(4.3 zk 2"2) × 102
(5"0 ± 1.0) × 103
L
1.7 × 102
3.7 × 102
Interhaloid complexes in aqueous solution
59
TABLE I.---EFFECT OF CONCENTRATION OF HYDROCHLORIC ACID ON rile REDOX POFENT1AL OF A SOLUTIONOF BROMINECHLORIDE IN HYDROCHLORIC ACID Redox potentials BrCI (mole)
HCI (mole)
0"005
0"05
TABLE 2 . - - E F F E C T
measured (mY)
calculated (mY)
0"2 0"5 I'0 1.5 2'0 2'5 3"0 4"0 5"0
1323 1290 1268 1237 1230 1209 1195 1170 1156
1317 1300 1263 1239 1219 1206 1191 1175 1154
0.6 1'0 1'5 2"0 2"5 3'0
1356 1336 1309 1283 1267 1253
1361 1332 1308 1288 1276 1260
OF CONCENTRATION OF HYDROCHLORIC ACID ON THE REDOX
POTENTIAL OF A SOLUTION OF IODINE CHLORIDE IN HYDROCHLORIC ACID
Redox potentials ICI (mole)
0.01
0'02
HCI (mole)
0'2 0"5 1"0 I '5 2'0 2"5 3"0
0"2 0"5 I'0 1"5 2"0 2"5 3"0
i
mesaured (mY)
calculated (mV)
1060 1040 1000 979 962 948 935
1064 1044 1005 978 957 944 931
1080 1050 1018 997 984 970 956
1084 1063 1024 997 977 963 953
60
E. PuNc_,oR, K. BUROEtXand E. Scm.~..Eg:
TABLE 3.--E~ECT
o F CONCENTRATION OF BROMIDE IONS ON ~
REDOX POTENTIAL
OF A SOLUTION OF IODINE BROMIDE CONTAINING BROMIDE IONS I
Rcdox potentials IBr (mole)
0"01
• 0"02
Br(mole)
measured (mV)
calculated (mV)
0-5 1.0 1"5 2"0 2'5 3"0
860 830 815 805 792 781
857 827 810 796 788 780
0"2 0"5 1-0 1"5 2.0 2-5 3"0
925 880 850 832 823 809 801
921 881 852 834 820 811 804
TABLE 4 . - - E F F E C T OF CONCENTRATION OF IODIDE IONS ON THE REDOX POTENTIAL OF A SOLUTION OF IODINE CONTAINING IODIDE IONS
Rcdox potentials Is
I-
(mole)
(mole)
measured (mY)
calculated (mV)
0"01
0"5 l'O 1"5 2"0 2"5
490 468 459 447 44O
479 465 456 449 444
0.02
0.5 1.0 1"5 2"0 2-5
503 481 469 460 456
497 482 473 466 462
lnterhaloitl complexes in aqueous solution
61
Since our calculations for both iodine chloride and bromide gave values of K and n differing from those given in the literature, it seemed necessary to verify the validity of our experimental method. Investigation of the tri-iodide complex seemed suitable for this purpose, the co-ordination number and stability constant having been calculated in a perfectly reliable manner. ~41 From redox potential measurements, followed by the use of equation (3), these parameters were calculated for the tri-iodide complex; the co-ordination number was found to be two, and the stability constant (3.2-q- 1.3) ~'< 102, (Table 4) in fair agreement with the published figure of 7 × 102. The correctness of the assumption involved in our method is thus demonstrated. '~ G. Jo~Es a n d B. B.
KAPLAN,J. Amer. Chem. Soc. 50, 1845 (1928).