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Spectrophotometric and potentiometric studies on the composition of iron complexes of potassium molybdo- and molybdicyanide
SPECTR~PHOT~M~TRIC AND P~TENTI~M~TRIC STUDIES ON THE COMPOSITION OF IRON COMPLEXES OF POTASSIUM MOLYBDO- AND MOLYBDICYANIDE WAHID U. MALIK and S. IFTIKX%R ALI ClxmkalLaboratories, Aligarh Mu&n ~niye~ity, Aligarh, India (Received 14 Mzrck 1961, Accepted 30 May 1961) S~-IronIlx forms a soIuble complex with potassium moly~ocy~ide. An inst~t~eous blue colour is produced when ferric chloride is added to a solution of potassium maly~ocy~ide. The composition of the complex was determined by the method of eontinuaus variation and by the slope ratio method, which indicated the existence of the complex KFelrlMolv(CN),. The formation constant of the complex was found to be 1.37 x LO*and the free energy of formation works out to be - 5.57 Kcals at 20”. Ferric sulphate also gives a blue complex with potassium molybdicyanide. Its composition has been studied potentiometrically using the MOM*+ Mo(CN),~- + ecouple at a bright platinum electrode; it was found to be K FeilMoV (CN),. SOLUBLE complexes of heavy metal molybdocyanides are not reported in the literature. We, for the first time, obtained a Crxrr complex of this type by the interaction of chromiumrlf chloride and potassium moly~o~a~de, and assigned the formula KCrJIrMorv (CN), to the red compound X,2thus obtained. Other soluble complexes of this series were obtained by the interaction of Ferrr with potassium molybdocyanide and Fen with potassium molybdocyan~de. This communication deals with the composition of these complexes as indicated by physi&o-chemical methods of examination. Preliminary experiments on the interaction of ironrrr chloride and ironrr sulphate with potassium molybdocyanide and potassium molybdicyanide respectively gave evidence of similar type of colour changes. Thus it was observed that with excess of ironrrf or ironrr ions, a blue solution was obtained, but with excess af molybdo- or rno~yb~~~de a greenish-blue solution is obtained. Since the reaction appeared to be anafogous to that observed in the formation of “iron blues”, it was thought worthwhile to elucidate the nature and composition of the complex or complexes formed, using various phys~co-che~cal methods. The composition of the complex formed by the interaction of ferric chloride and potassium mo~ybdocya~ide could not be studied satisfactorily by conductomet~~~ potentiometer or amperometric titration methods. ~nteres~gly enough this complex was found to be reducible at the dropping mercury electrode, giving a typical pohrographic wave, but the polarographic method3 could not be successfully employed either in view of the fact that the complex became unstable on coming in contact with the mercury drops (a slight turbidity, followed by precipitation in the polarographic cell, was observed). The spectrophotometric method was, however, found useful in determining the composition and stability of the complex. This method was also tried for investigating the complex formed by the interaction of ferric sulphate and potassium molyb~cya~de, but reliable information could not be obtained since K~~o(C~s was appreciably decomposed by the light from the tungsten lamp of the
WAHID U. MALJK and S. Irrncm~
738
ALS
Several electrometric methods were then tried and it was found sp~ctro~~otometer. that very satisfactory results could be realised by a potent~omet~c method, using a MO(CNj$-- + Mo(CN)s3- + e- couple at a bright platinum electrode. EXPERIMENTAL
Reagents Ferric chloride (AnalaR) and ferrous suEphnte(recrystallised three times) were dissolved in doubly distilled water and the strengths of the solutions were determined by the usual methods. Potassium molybdocyanide was prepared by the method recommended by Fieser’ and the strength of the solution was determined by titrating ~tentiome~i~y against KM& solution. Kolthoff’s methods was used for obtaining ~~~ass~~rnrno~~~~c~~n~~. TIE strength of the solution was also determined potenti~metri~ly using potassium ferrocyanide. ~~~~a~~ A Reckman DU s~ct~pbotometer with l-cm Corex cells and a tun~ten was used for optical density measurements.
lamp as the light source
PROCEDURES In order to determine the number of complexes formed by the interaction of ferric chloride and potassium molybdocyanide, the method of Vosburgh and Coopers was followed. The reactants were mixed in different, proportions and the optical density was measured at different wavelengths ranging from 360 rnp to 1000 rnp. All the mixtures gave a maximum at 840 rnfi showing thereby the formation of only one complex (Fig. 1).
,
00
Ro. I.-FeCiS
and &Mo(CN), mixed in the proportions (1) 2: 1, (2) 1: 1, and (3) 1:2. Concentration of the reactants, 3-3 X 1O-sM.
Job’s method? of continuous vocation was followed for ~te~~~ng the composition of ferric molybdocyanide. Solutions of ferric chIoride and potassium moiybdo~~de of three different concentrations, viz., 4 x 10+&f, 2 x 10-aM and 1.25 x lOaM, were mixed according to the method of continuous variation, and O.D. was measured at 840 rnfi after 10 min of mixing. Ferric chloride and potassium molybdocyanide at these concentrations had a negligibly small absorption at 84Omg. The optical densities of the solutions were plotted against the ratio [Fe*+]/[Fe*+] t [MO&N),*-1. The results are depicted in Fig. 2. The results were further co&rmed by the slope ratio method.* Two sets of event were performed. In the first set the ~n~ntmti~ of potassium ~Iy~ocyanide was kept constant and that of ferric chloride was varied; in the other set the mixing was done in the reverse order. The results are given in Fig. 3. The slope was determined over the straight lime portion of the curve.
Composition
of iron complexes of potassium molybdo- and molybdicyanide
Fro. 2. The ratio FeYFe*+ -tMo(CN),~. Concentrationsof thereactants, (1) 4 x lo-skf, (2) 2 X 104&f and (3) 1.25 x IO-SM.
1*60-
1*40-
FIG. 3. Slope ratio method. (1) Concentration of K,Mo(CN)s constant, 1.66 X IOWikf. Concentration of FeC& varied, 2 X IO-*M. (2) Concentration of FeCl, constant, I.66 x 10-sM. Concentration of K,Mo(CN), varied, 2 X IO+M.
z.. I.20 t* 2 0 r*oo0 .o 2 0.60-
0.60 -
I
0
I
0.1
0.2
,
0.3
O-4
1
0.5
0.6
0.7
WAHXDU. MALIK and S. IFTIKHAR ALI
740
Potentiometric titrations between ferrous sulphate and potassium molybdicyanide were carried out using a Tinsley Vernier potentiometer (type 33873); the indicator electrode [Mo(CN),*- + Mo(CN),~- + e- coupIe] was obtained by dipping a bright platinum electrode in the solution of potassium molybdicyanide containing a little potassium molybdocyanide. The titrations at three different concentrations of molybdicyanide solution, viz., 0.03&f, 0.02M and 0~02M, were performed for determining the composition of ferrous molybdicyanide. In the cell fitted with the indicator electrode and the saturated calomel electrode, 10 ml of the molybdicyanide solution were taken and ferrous sulphate (O.lM) was added from the burette. The curves are shown in Pig.4. Reverse titrations with ferrous sulphate in the cell were not successful. The experiments were performed in a dark room and the vessels were wrapped with black paper to avoid the decomposition of potassium molybdicyanide by light.
0.70
-
O-65 -
0.60 -
0.55 > o-50P 2 xl o-45 -
0.40 -
0.35 -
o-30 -
0.25 0.201 0.5
f
I.0
8
I.5
I 2.0 FeSO,,
/
2.5 mL
3.0
t
3.5
FIG. 4.-Determination of end-point in potentiometric titration. IO ml of K,Mo(CN), concentration (1) 0*03M, (2) @025M, and (3) 0.02OM versus O.lM FeSO,.
of
DISCUSSION
Several authors, including Davidson: have extensively investigated the nature of the reaction between ferrous chloride and potassium ferrocyanide and the composition of the complex. When irorP ions are added to an equimolecular quantity of ferrocyanide ions, a redox equilibrium sets in with the result that the iron’u ions are almost completely reduced to iro@ ions by the ferrocyanide ions, which, in turn, are simultaneously oxidised to ferricyanide ions. The possibility of such a redox
Composition of iron complexes of potassium molybdo- and molybdicyanide
741
equilibrium, existing in the case of the ferric chloride-potassium molybdocyanide reaction appears to be remote, since the equilibrium constant value,
K298= ~Fe~~l[Mo~C~~*-l = X0.4, {Fe2*]~Mo(C~s3-]
(Fe2+ + Fe3f + e, E, = -0.76; + Mo(CN)s3- + e, E, = -0.82)
MO(W),*-
is quite low when compared to the ferric chloride-potassium ferrocyanide reaction (ly298 = 1.215 x 103. Thus on theoretical grounds it could be inferred that the complexes formed by the interaction of ferric chloride and potassium molybdocyanide, and between ferrous sulphate and potassium molybdicyanide, could not be one and the same and unlike “iron blues” would have different compositions. The combining ratio of 1 :l for Fe 3+ to Mo(CN)s*- is obtained both by the continuous variation and slope ratio methods (Figs. 2 and 3) for ihe interaction of ferric chloride and potassium molybdi~ya~de. The composition of the complex would, therefore, be represented by the formula KFelllMo’V(CN)s, according to the reaction, FeCl, + &Mo(CN),
-+ KFelIIMorv(CN),
+ 3 KCl.
The formation constant of the complex was determined by the method followed by Mukherji and Deyl* for the reaction mA + nB s A~BPI, where m/n=lorm=lorm=~=l The formation constant is X
K=
(a - x)(b -
x) ’
where x is the concentration of the complex, and a and b are the initial concentration of ferric chloride and potassium molybdocyanide respectively. Taking two concentrations a, and a2 and b, and b, of the reactants giving the same optical density. (that is the same value of x), K= or
X (01
-
X)(b,
-
x)
=
(a2
-
x;&,
-
x)
a&, - a&,
x=(a, + b,) -
(a2 + bd
’
Knowing the value of x from the above equation, K is calculated from different values of ta and b. Taking two concentrations of Fe3+ and Mo(CN),*- at the optical density 0.20 (Fig. 2), and the value of x from the above equation as 2.19 x fO-4&f, the value of K was found to be 1.37 x 104. The free energy of formation of the complex was calculated with the help of the expression AF” = -RTlnk, where AF” is the standard free energy change, R, the gas constant and T, the absolute temperature. The free energy of formation of ferric molybdocyanide complex works out to be -5-75 kcals. at 20”. Potentiometric titrations carried out at different concentrations of potassium molybdicyanide with ferrous sulphate give the combining ratio of Fe2f to Mo(CN)s3-
142
WAHID
U. MALIK and S. IFTIKHAR ALI
as 1 :l, indicating the formation of the complex KFe”Mo” reaction, FeSO, + KsMo(CN)s + KFeI’Mo”(CN),
(CN),, according to the + K&l,.
This method, in addition to providing information regarding the composition of the ironrr molybdicyanide complex, has demonstrated the utility of a new electrode (the Mo(CN),~- + Mo(CN)s3- + e- system at bright platinum) for investigating the reaction between heavy metal ions and potassium molybdo- and molybdicyanide. This electrode could be used successfully for studying the precipitation reaction between potassium molybdocyanide and NP, Corr, or Curr. Zusanuneafassung-Auf Zugabe von Eisen(III)-ionen zu einer Losung von Kalium-Molybdocyanid bildet sich augenblicklich ein blauer, wasserltislicher Komplex. Die Zusammensetzung des Komplexes wurde auf grund eines Job-Diagrammes und nach der Methode der Neigungsverhaltnisse als folgende ermittelt: KFe(III)Mo(IV)(CN),. Die Bildungskonstante betragt 1.37 x 104. Die freie Bildungsenergie wurde zu -5.57 Cal bei 20°C berechnet. FerrisuIfat gibt such mit Molybdicyanid einen blauen Komplex. Seine Zusammensetzung wurde durch Potentiahnessungen an Gleichgewicht Mo(CN)-~ + Mo(CN),-” + e- an einer blanken Platinelektrode ermittelt, und ist KFe(III)Mo(V) (CN),. R&stun&Par addition de fer(II1) a une solution de molybdo-cyanure de potassium, il se forme instantanement un complexe bleu soluble. La composition du complexe a Cte determin6e par la methode des variations continues et par la methode du rapport des pentes; ces methodes montrent La constante de formation du complexe est 1,37 . lo4 l’existence du complexe KFe~tlr’Mouv)(CN),. et l’energie libre de formation calcul&e est -5,57 kcal. a 20”. Le sulfate de fer(I1) donne aussi un complexe bleu avec le molybdo-cyanure de potassium. Sa composition a CtC Ctudi6e par potentiometric a une electrode de platine poli, utilisant le systeme oxydo-reducteur: Mo(CN),~-
+ Mo(CN),~- + e
Les auteurs ont trouve la formule suivante pour le complexe:
KFe(tl’Mo’V)(CN),.
REFERENCES 1 W. U. Malik and S. Iftikhar Ali, Nuturwiss., 1959, 20, 579. 2 Idem, J. Inorg. Nuclear Chem., in press. a James J. Lingane, Chem. Revs., 1941, 29, 1. 4 Louis F. Fieser, J. Amer. Chem. Sot., 1930, 52, 5226. 5 I. M. Kolthoff and W. J. Tomsicek, J. Phys. Chem., 1936, 40,247. 6 W. C. Vosburgh and G. R. Cooper, J. Amer. Chem. Sot., 1941,63,437. ’ P. Job, Ann. Chim., 1928, 9, (lo), 113. BE. Harvey and D. L. Manning, J. Amer. Chem. Sot., 1950,72,4488. 9 D. Davidson, J. Chem. Educ., 1937, 14, 238. lo A. K. Mukherji and A. K. Dey, Proc. Nat. Acad. Sci. India, 1957, 26, 20.
WAHID U. MALJK and S. Irrncm~
738
ALS
Several electrometric methods were then tried and it was found sp~ctro~~otometer. that very satisfactory results could be realised by a potent~omet~c method, using a MO(CNj$-- + Mo(CN)s3- + e- couple at a bright platinum electrode. EXPERIMENTAL
Reagents Ferric chloride (AnalaR) and ferrous suEphnte(recrystallised three times) were dissolved in doubly distilled water and the strengths of the solutions were determined by the usual methods. Potassium molybdocyanide was prepared by the method recommended by Fieser’ and the strength of the solution was determined by titrating ~tentiome~i~y against KM& solution. Kolthoff’s methods was used for obtaining ~~~ass~~rnrno~~~~c~~n~~. TIE strength of the solution was also determined potenti~metri~ly using potassium ferrocyanide. ~~~~a~~ A Reckman DU s~ct~pbotometer with l-cm Corex cells and a tun~ten was used for optical density measurements.
lamp as the light source
PROCEDURES In order to determine the number of complexes formed by the interaction of ferric chloride and potassium molybdocyanide, the method of Vosburgh and Coopers was followed. The reactants were mixed in different, proportions and the optical density was measured at different wavelengths ranging from 360 rnp to 1000 rnp. All the mixtures gave a maximum at 840 rnfi showing thereby the formation of only one complex (Fig. 1).
,
00
Ro. I.-FeCiS
and &Mo(CN), mixed in the proportions (1) 2: 1, (2) 1: 1, and (3) 1:2. Concentration of the reactants, 3-3 X 1O-sM.
Job’s method? of continuous vocation was followed for ~te~~~ng the composition of ferric molybdocyanide. Solutions of ferric chIoride and potassium moiybdo~~de of three different concentrations, viz., 4 x 10+&f, 2 x 10-aM and 1.25 x lOaM, were mixed according to the method of continuous variation, and O.D. was measured at 840 rnfi after 10 min of mixing. Ferric chloride and potassium molybdocyanide at these concentrations had a negligibly small absorption at 84Omg. The optical densities of the solutions were plotted against the ratio [Fe*+]/[Fe*+] t [MO&N),*-1. The results are depicted in Fig. 2. The results were further co&rmed by the slope ratio method.* Two sets of event were performed. In the first set the ~n~ntmti~ of potassium ~Iy~ocyanide was kept constant and that of ferric chloride was varied; in the other set the mixing was done in the reverse order. The results are given in Fig. 3. The slope was determined over the straight lime portion of the curve.
Composition
of iron complexes of potassium molybdo- and molybdicyanide
Fro. 2. The ratio FeYFe*+ -tMo(CN),~. Concentrationsof thereactants, (1) 4 x lo-skf, (2) 2 X 104&f and (3) 1.25 x IO-SM.
1*60-
1*40-
FIG. 3. Slope ratio method. (1) Concentration of K,Mo(CN)s constant, 1.66 X IOWikf. Concentration of FeC& varied, 2 X IO-*M. (2) Concentration of FeCl, constant, I.66 x 10-sM. Concentration of K,Mo(CN), varied, 2 X IO+M.
z.. I.20 t* 2 0 r*oo0 .o 2 0.60-
0.60 -
I
0
I
0.1
0.2
,
0.3
O-4
1
0.5
0.6
0.7
WAHXDU. MALIK and S. IFTIKHAR ALI
740
Potentiometric titrations between ferrous sulphate and potassium molybdicyanide were carried out using a Tinsley Vernier potentiometer (type 33873); the indicator electrode [Mo(CN),*- + Mo(CN),~- + e- coupIe] was obtained by dipping a bright platinum electrode in the solution of potassium molybdicyanide containing a little potassium molybdocyanide. The titrations at three different concentrations of molybdicyanide solution, viz., 0.03&f, 0.02M and 0~02M, were performed for determining the composition of ferrous molybdicyanide. In the cell fitted with the indicator electrode and the saturated calomel electrode, 10 ml of the molybdicyanide solution were taken and ferrous sulphate (O.lM) was added from the burette. The curves are shown in Pig.4. Reverse titrations with ferrous sulphate in the cell were not successful. The experiments were performed in a dark room and the vessels were wrapped with black paper to avoid the decomposition of potassium molybdicyanide by light.
0.70
-
O-65 -
0.60 -
0.55 > o-50P 2 xl o-45 -
0.40 -
0.35 -
o-30 -
0.25 0.201 0.5
f
I.0
8
I.5
I 2.0 FeSO,,
/
2.5 mL
3.0
t
3.5
FIG. 4.-Determination of end-point in potentiometric titration. IO ml of K,Mo(CN), concentration (1) 0*03M, (2) @025M, and (3) 0.02OM versus O.lM FeSO,.
of
DISCUSSION
Several authors, including Davidson: have extensively investigated the nature of the reaction between ferrous chloride and potassium ferrocyanide and the composition of the complex. When irorP ions are added to an equimolecular quantity of ferrocyanide ions, a redox equilibrium sets in with the result that the iron’u ions are almost completely reduced to iro@ ions by the ferrocyanide ions, which, in turn, are simultaneously oxidised to ferricyanide ions. The possibility of such a redox
Composition of iron complexes of potassium molybdo- and molybdicyanide
741
equilibrium, existing in the case of the ferric chloride-potassium molybdocyanide reaction appears to be remote, since the equilibrium constant value,
K298= ~Fe~~l[Mo~C~~*-l = X0.4, {Fe2*]~Mo(C~s3-]
(Fe2+ + Fe3f + e, E, = -0.76; + Mo(CN)s3- + e, E, = -0.82)
MO(W),*-
is quite low when compared to the ferric chloride-potassium ferrocyanide reaction (ly298 = 1.215 x 103. Thus on theoretical grounds it could be inferred that the complexes formed by the interaction of ferric chloride and potassium molybdocyanide, and between ferrous sulphate and potassium molybdicyanide, could not be one and the same and unlike “iron blues” would have different compositions. The combining ratio of 1 :l for Fe 3+ to Mo(CN)s*- is obtained both by the continuous variation and slope ratio methods (Figs. 2 and 3) for ihe interaction of ferric chloride and potassium molybdi~ya~de. The composition of the complex would, therefore, be represented by the formula KFelllMo’V(CN)s, according to the reaction, FeCl, + &Mo(CN),
-+ KFelIIMorv(CN),
+ 3 KCl.
The formation constant of the complex was determined by the method followed by Mukherji and Deyl* for the reaction mA + nB s A~BPI, where m/n=lorm=lorm=~=l The formation constant is X
K=
(a - x)(b -
x) ’
where x is the concentration of the complex, and a and b are the initial concentration of ferric chloride and potassium molybdocyanide respectively. Taking two concentrations a, and a2 and b, and b, of the reactants giving the same optical density. (that is the same value of x), K= or
X (01
-
X)(b,
-
x)
=
(a2
-
x;&,
-
x)
a&, - a&,
x=(a, + b,) -
(a2 + bd
’
Knowing the value of x from the above equation, K is calculated from different values of ta and b. Taking two concentrations of Fe3+ and Mo(CN),*- at the optical density 0.20 (Fig. 2), and the value of x from the above equation as 2.19 x fO-4&f, the value of K was found to be 1.37 x 104. The free energy of formation of the complex was calculated with the help of the expression AF” = -RTlnk, where AF” is the standard free energy change, R, the gas constant and T, the absolute temperature. The free energy of formation of ferric molybdocyanide complex works out to be -5-75 kcals. at 20”. Potentiometric titrations carried out at different concentrations of potassium molybdicyanide with ferrous sulphate give the combining ratio of Fe2f to Mo(CN)s3-
142
WAHID
U. MALIK and S. IFTIKHAR ALI
as 1 :l, indicating the formation of the complex KFe”Mo” reaction, FeSO, + KsMo(CN)s + KFeI’Mo”(CN),
(CN),, according to the + K&l,.
This method, in addition to providing information regarding the composition of the ironrr molybdicyanide complex, has demonstrated the utility of a new electrode (the Mo(CN),~- + Mo(CN)s3- + e- system at bright platinum) for investigating the reaction between heavy metal ions and potassium molybdo- and molybdicyanide. This electrode could be used successfully for studying the precipitation reaction between potassium molybdocyanide and NP, Corr, or Curr. Zusanuneafassung-Auf Zugabe von Eisen(III)-ionen zu einer Losung von Kalium-Molybdocyanid bildet sich augenblicklich ein blauer, wasserltislicher Komplex. Die Zusammensetzung des Komplexes wurde auf grund eines Job-Diagrammes und nach der Methode der Neigungsverhaltnisse als folgende ermittelt: KFe(III)Mo(IV)(CN),. Die Bildungskonstante betragt 1.37 x 104. Die freie Bildungsenergie wurde zu -5.57 Cal bei 20°C berechnet. FerrisuIfat gibt such mit Molybdicyanid einen blauen Komplex. Seine Zusammensetzung wurde durch Potentiahnessungen an Gleichgewicht Mo(CN)-~ + Mo(CN),-” + e- an einer blanken Platinelektrode ermittelt, und ist KFe(III)Mo(V) (CN),. R&stun&Par addition de fer(II1) a une solution de molybdo-cyanure de potassium, il se forme instantanement un complexe bleu soluble. La composition du complexe a Cte determin6e par la methode des variations continues et par la methode du rapport des pentes; ces methodes montrent La constante de formation du complexe est 1,37 . lo4 l’existence du complexe KFe~tlr’Mouv)(CN),. et l’energie libre de formation calcul&e est -5,57 kcal. a 20”. Le sulfate de fer(I1) donne aussi un complexe bleu avec le molybdo-cyanure de potassium. Sa composition a CtC Ctudi6e par potentiometric a une electrode de platine poli, utilisant le systeme oxydo-reducteur: Mo(CN),~-
+ Mo(CN),~- + e
Les auteurs ont trouve la formule suivante pour le complexe:
KFe(tl’Mo’V)(CN),.
REFERENCES 1 W. U. Malik and S. Iftikhar Ali, Nuturwiss., 1959, 20, 579. 2 Idem, J. Inorg. Nuclear Chem., in press. a James J. Lingane, Chem. Revs., 1941, 29, 1. 4 Louis F. Fieser, J. Amer. Chem. Sot., 1930, 52, 5226. 5 I. M. Kolthoff and W. J. Tomsicek, J. Phys. Chem., 1936, 40,247. 6 W. C. Vosburgh and G. R. Cooper, J. Amer. Chem. Sot., 1941,63,437. ’ P. Job, Ann. Chim., 1928, 9, (lo), 113. BE. Harvey and D. L. Manning, J. Amer. Chem. Sot., 1950,72,4488. 9 D. Davidson, J. Chem. Educ., 1937, 14, 238. lo A. K. Mukherji and A. K. Dey, Proc. Nat. Acad. Sci. India, 1957, 26, 20.