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Formation of a silicato-iron(III) complex in dilute aqueous solution
J. I n o r g . N u c l , C h e m ,
1965, Vol. 27. pp. 237 to 239,
Pergamon
Press Ltd.
Printed
in Northern
Ireland
NOTES
Formltion of a silicato-iron(lll) complex in dilute aqueous solution (Received 16 March 1964; in revisedform 22 June 1964) Qu^~.ITA'nvEevidence for the chemical interaction of silicic acid with iron(III) in dilute aqueous solution has been reported, cx~but neither information regarding the nature of the interaction nor of the chemical species formed is available. Uv. absorption data in dilute acidic aqueous solutions indicate that Fe *+ forms relatively stable FeSiO(OH),2+ complexes with Si(OH),. EXPERIMENTAL The experimental procedures and methods for evaluation of the data for the present investigations were similar in principle to those reported earlier c*~for study of the reaction of Fe(III)with orthophosphate• All experiments were carried out at a temperature of 25° ± 0.5°C in CIO,- solutions of ionic strength,/~ = 0-I. The experimental conditions, [H+] > 10-a's M and [Fe]m ~ 10-' M, were so chosen that only the first hydrolysis equilibrium of Fe(III) had to be considered. ~s~ Stock solutions of iron(III) were prepared by dissolving appropriate quantities of hydrated ferric • • . ~ . • , perchlorate (G. Fred. Srmth) m dilute HCIO4. The concentration of lron(III) was deternuned after dilution to less than 10-' M and adjustment to [H +] > 10-*'s from absorbance measurements at a wave length of 272 m/~, the isobestic point of the Fe~--FeOH*+ equilibrium. Stock solutions of monomeric orthosilicic acid, Si(OH)~, were prepared according to a method proposed by At~XNDER,~ by passing solutions of sodium metasilicate (Na~SiO. • 9H~O) through a 30 cm column of cation exchange resin (Dowex 50W-X8) in the hydrogen ion form. In some of the stock solutions so prepared the Si(OH), concentration (ranging up to 5 x 10-~ M) exceeded the solubility of amorphous SiOs (2 x 10-* M). cs~ Since such oversaturated solutions undergo slow polymerization,~6~a freshly prepared solution was used for each experiment. The silica in all solutions used was found to react completely with molybdic acid within 1-2 min, thus indicating that all the silicic acid existed as the monomeric species.~T~ The concentration of Si(OH), was computed from the weight of sodium metasilicate used. Absorbance measurements were made with a Beckman Model DU spectrophotometer using cells of l0 cm light path. After preparation, each solution was stored for 1 hr in a constant temperature bath prior to measurements of absorbance. Hydrogen ion concentration was computed from the amount of standard HCIO, added to the solution. Hydrogen ion contribution due to hydrolysis of Fe s+ and complex formation was negligible. A N A L Y S I S OF S P E C T R O P H O T O M E T R I C D A T A The addition of Si(OH), to acid ferric perchlorate solutions increases light absorption within the wavelength range 240-300 m/~. Pure silicic acid solutions do not absorb at these wavelengths. For evaluation of the extent of interaction of iron(II1)and silicic acid in the mixed solutions, absorption measurements at a wavelength of 272 rn~ for different mixtures were analysed. This wavelength corresponds to the isobestic point of the FeS+-FeOH *+ equilibrium. The data were treated by using ~t~F. HAZEL,R. U. ScHocK, JR. and M. GORDON,J. Amer. Chem. Soc. 71, 2256 (1949). ~a~H. GALAL-GORCHEVand W. STUI~, J. Inorg. Nuci. Chem. 25, 567 (1963). ~a~R. M. MILaURN,J. Amer. Chem. Soc. 79, 537 (1957). c,~ G. B. ALEXANDER,J. Amer. Chem. Soc. 75, 2887 (1953). ~s~G. LAG~nsTR6M,Acra chem. scand. 13, 722 0959). ~s~K. Goro, J. phys. Chem. 60, 1007 0956). ~ G. B. A~XA~DER,W• M. H~roN and R. K. I~Elt, J. phys. Chem. 58, 453 0954). 237
238
Notes
a modification(') of a method described by NEWTONand ARCAND.(6) The experimental results may be interpreted by considering the following equilibria: Fe a~ + Si(OH)~ = FeH,SiO4 f . - ' + + (4 -- n)H+; Ql Fe *+ + H20 = FeOH ~+ + H+; QH
(1) (2)
The Q's are equilibrium quotients; for example, Qu = [FeOH2~][H+]/[ Fe3 ~], where the brackets, [ ], signify concentration in moles/l. If the absorbance of a mixed solution of Fe(IIl) and Si(OH)4 is represented by A, the absorbance for the Fe(III) in the absence of Si(OH), by A0, and the hypothetical absorbance if all the Fe(i ll) were in the form of the silicato-iron(III) complex by A 1, then: A/1 = ere([Fes÷] + [FeOH2+]) + ereH,siot [FeH,SiO4 ¢"-t)+]
(3)
A,/1 = eveu.slo,[FeT]
(4)
Ao/1 = eFe[Fe~]
(5)
where eve and eve.,sio, are the molar absorbances at 272 m# o f Fe s+ (or FeOH *+) and FeH,SiO Ltn-1)+, respectively, and Fe~ is given by [Fe~] = [Fes+] + [FeOH s+] + [FeH~SiO, c"-1)+]
(6)
For direct application to experimental data Equations (1) through (6) are combined to: A
=
AI
- -
(.4 --Ao)[H+]~'-")[1 + Q~/[H+]] QxISi(OH),]
(7)
Thus a plot of A vs. (A -- Ae)/[Si(OH),] for a series of soluiions which have varying concentrations of Si(OH), but have the same formal concentration of Fe~ and H + should give a straight line, provided that within the concentration regions studied only one silicato complex species predominates. At and [H]t4"-")/Qx can be evaluated from the intercept and slope of such a plot. In applying Equation (7) to the experimental data the concentration of Si(OH)4 was set equal to the total concentration of dissolved silica, Si~, since within the concentration and pH range studied the ionization of Si(OH)( (K = 10-''5¢6)) was negligible, and, in addition, [FeH,SiO,] ~ (SIT). RESULTS Data for several series of experimental systems are plotted according to Equation (7) in Fig. I. Each line represents a fit of the data for a series of solutions of constant (H+). Good normalization of the data over a rather wide range of Si(OH), concentrations has been obtained for solutions with total iron(III) concentrations of less than 5 X 10-~ M; such a linear relationshipis expected for the formation of a complex that contains one mole ligand per mole of metal ion. re) (For a complex with z ligands per metal ion, the [Si(OH,] in the abscissa term would have to be raised to the z ta power in order to obtain normalization of the data.) Inconsistent results were obtained when ferric solutions more concentrated than 5 x 10-5 M were used. This is probably attributable to the formation of polynuclear complexes. Evaluation of n and Q: from the data given in Fig. 1 is made by rearranging Equation (l) in the logarithmic form log [[FeH.SiO.(--:)+]/[Fe*+][Si(OH).]] = log Qt - (4 - n) log[H +]
(8)
and by plotting the left side of this Equation vs. --log[H+]. The ordinate values in Fig. 1 (inset) were obtained by dividing the quantity [1 + Qa/[H+]] for each series of solutions by the slope of the corresponding trace shown in Fig. I. A value of Q , = 2.89 x 10-s (/z = 0.1, 25°C) (s) has been used in the calculations. From the logarithmic plot values of n = 3 and Q = 0.57 can be derived. Thus, the equilibrium relationships for the complex formation between Fe(III) and orthosilicic acid under the experimental conditions studied ([H +] < 10-s's and [Fe(III)] < 10-4"s) can be represented by Fe s+ + Si(OH), = FeSiO(OH), a+ + H+; Qx = 0.57 (25°C, # = 0.1) t,) T. W. Nv.wa'oN and M. G. ARCAND, J. Amer. Chem. Soc. 75¢ 2449 (1953).
(9)
Notes
239 <3
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350(1
~
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2,5
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I
1
\
iO-IJlO ;5000
3.0
-leg {H+1
I
9-
3
\
4,
5
.A-~
6
7
8
9
,
[Si(OH)4][FeT] {
x I0 "5
Fig. 1
TM
Using 10-.'s as the first concentration acidity constant of Si(OH), the following value of the concentration stability constant for the silicato-ferri complex is obtained: K = Qx/Ksl~om, = [FeSiO(OH)=S+]/[Fe*+][SiO(OH)s-] -~ 1"8 × 10' (25°C,/4 = 0"1)
(10)
It is interesting to compare such complex formation in homogeneous solution with ferric ion and general cation exchange reactions that are ob~rved with dlica gel; ct,loj this ion exchange has been interpreted as surface complex formation of the metal ion with the weakly acidic silanol (--SiOH) groups of the silica gel surface. Natural waters contain as much as 10-s M dis.qolved silica. The results reported here, confirm that iron(IIl) and silicates can interact chemically under certain conditions; such interactior~ have been pmtulated in connection with observations on limnological transformations of iron and silicaJ TM In evaluating the effect of silicates on coagulation by ferric iron ~lt~specific chemical interactions have to be considered in addition to those of the electric double layer. Acknowledgement--This work was supported in part by Public Health Service Grant WP 00043. Division of F~glneering and Applied Physics Harvard University Cambridge 38, Mass.
W . J . Wnleg, Jlt.* W. STUMM
* Present address: Department of Civil Engineering, University of Michigan, Ann Arbor, Michigan. c,~ j. STAWnm and R. W. M^Aa'U~,N,J. Coil. Sci. 18, 132 (1963). ,o~ D. L. DUOOER, J. H. ST~WrON, B. N. I u Y , B. L. McCom~eLL, W. W. Cuunmt,~3s and R. W. MAA'rMAN,J. phys. Chem. 68, 757 (1964). cxx~j. j. MOROANand W. S ~ , in Proe. 2nd Int. Conf. on Water Pollution Control, Tokyo 1964, Pergamon, London. ~ls~j. R. BAYLIS,J. Amer. Water Works Ass. 29, 1355 (1937).
1965, Vol. 27. pp. 237 to 239,
Pergamon
Press Ltd.
Printed
in Northern
Ireland
NOTES
Formltion of a silicato-iron(lll) complex in dilute aqueous solution (Received 16 March 1964; in revisedform 22 June 1964) Qu^~.ITA'nvEevidence for the chemical interaction of silicic acid with iron(III) in dilute aqueous solution has been reported, cx~but neither information regarding the nature of the interaction nor of the chemical species formed is available. Uv. absorption data in dilute acidic aqueous solutions indicate that Fe *+ forms relatively stable FeSiO(OH),2+ complexes with Si(OH),. EXPERIMENTAL The experimental procedures and methods for evaluation of the data for the present investigations were similar in principle to those reported earlier c*~for study of the reaction of Fe(III)with orthophosphate• All experiments were carried out at a temperature of 25° ± 0.5°C in CIO,- solutions of ionic strength,/~ = 0-I. The experimental conditions, [H+] > 10-a's M and [Fe]m ~ 10-' M, were so chosen that only the first hydrolysis equilibrium of Fe(III) had to be considered. ~s~ Stock solutions of iron(III) were prepared by dissolving appropriate quantities of hydrated ferric • • . ~ . • , perchlorate (G. Fred. Srmth) m dilute HCIO4. The concentration of lron(III) was deternuned after dilution to less than 10-' M and adjustment to [H +] > 10-*'s from absorbance measurements at a wave length of 272 m/~, the isobestic point of the Fe~--FeOH*+ equilibrium. Stock solutions of monomeric orthosilicic acid, Si(OH)~, were prepared according to a method proposed by At~XNDER,~ by passing solutions of sodium metasilicate (Na~SiO. • 9H~O) through a 30 cm column of cation exchange resin (Dowex 50W-X8) in the hydrogen ion form. In some of the stock solutions so prepared the Si(OH), concentration (ranging up to 5 x 10-~ M) exceeded the solubility of amorphous SiOs (2 x 10-* M). cs~ Since such oversaturated solutions undergo slow polymerization,~6~a freshly prepared solution was used for each experiment. The silica in all solutions used was found to react completely with molybdic acid within 1-2 min, thus indicating that all the silicic acid existed as the monomeric species.~T~ The concentration of Si(OH), was computed from the weight of sodium metasilicate used. Absorbance measurements were made with a Beckman Model DU spectrophotometer using cells of l0 cm light path. After preparation, each solution was stored for 1 hr in a constant temperature bath prior to measurements of absorbance. Hydrogen ion concentration was computed from the amount of standard HCIO, added to the solution. Hydrogen ion contribution due to hydrolysis of Fe s+ and complex formation was negligible. A N A L Y S I S OF S P E C T R O P H O T O M E T R I C D A T A The addition of Si(OH), to acid ferric perchlorate solutions increases light absorption within the wavelength range 240-300 m/~. Pure silicic acid solutions do not absorb at these wavelengths. For evaluation of the extent of interaction of iron(II1)and silicic acid in the mixed solutions, absorption measurements at a wavelength of 272 rn~ for different mixtures were analysed. This wavelength corresponds to the isobestic point of the FeS+-FeOH *+ equilibrium. The data were treated by using ~t~F. HAZEL,R. U. ScHocK, JR. and M. GORDON,J. Amer. Chem. Soc. 71, 2256 (1949). ~a~H. GALAL-GORCHEVand W. STUI~, J. Inorg. Nuci. Chem. 25, 567 (1963). ~a~R. M. MILaURN,J. Amer. Chem. Soc. 79, 537 (1957). c,~ G. B. ALEXANDER,J. Amer. Chem. Soc. 75, 2887 (1953). ~s~G. LAG~nsTR6M,Acra chem. scand. 13, 722 0959). ~s~K. Goro, J. phys. Chem. 60, 1007 0956). ~ G. B. A~XA~DER,W• M. H~roN and R. K. I~Elt, J. phys. Chem. 58, 453 0954). 237
238
Notes
a modification(') of a method described by NEWTONand ARCAND.(6) The experimental results may be interpreted by considering the following equilibria: Fe a~ + Si(OH)~ = FeH,SiO4 f . - ' + + (4 -- n)H+; Ql Fe *+ + H20 = FeOH ~+ + H+; QH
(1) (2)
The Q's are equilibrium quotients; for example, Qu = [FeOH2~][H+]/[ Fe3 ~], where the brackets, [ ], signify concentration in moles/l. If the absorbance of a mixed solution of Fe(IIl) and Si(OH)4 is represented by A, the absorbance for the Fe(III) in the absence of Si(OH), by A0, and the hypothetical absorbance if all the Fe(i ll) were in the form of the silicato-iron(III) complex by A 1, then: A/1 = ere([Fes÷] + [FeOH2+]) + ereH,siot [FeH,SiO4 ¢"-t)+]
(3)
A,/1 = eveu.slo,[FeT]
(4)
Ao/1 = eFe[Fe~]
(5)
where eve and eve.,sio, are the molar absorbances at 272 m# o f Fe s+ (or FeOH *+) and FeH,SiO Ltn-1)+, respectively, and Fe~ is given by [Fe~] = [Fes+] + [FeOH s+] + [FeH~SiO, c"-1)+]
(6)
For direct application to experimental data Equations (1) through (6) are combined to: A
=
AI
- -
(.4 --Ao)[H+]~'-")[1 + Q~/[H+]] QxISi(OH),]
(7)
Thus a plot of A vs. (A -- Ae)/[Si(OH),] for a series of soluiions which have varying concentrations of Si(OH), but have the same formal concentration of Fe~ and H + should give a straight line, provided that within the concentration regions studied only one silicato complex species predominates. At and [H]t4"-")/Qx can be evaluated from the intercept and slope of such a plot. In applying Equation (7) to the experimental data the concentration of Si(OH)4 was set equal to the total concentration of dissolved silica, Si~, since within the concentration and pH range studied the ionization of Si(OH)( (K = 10-''5¢6)) was negligible, and, in addition, [FeH,SiO,] ~ (SIT). RESULTS Data for several series of experimental systems are plotted according to Equation (7) in Fig. I. Each line represents a fit of the data for a series of solutions of constant (H+). Good normalization of the data over a rather wide range of Si(OH), concentrations has been obtained for solutions with total iron(III) concentrations of less than 5 X 10-~ M; such a linear relationshipis expected for the formation of a complex that contains one mole ligand per mole of metal ion. re) (For a complex with z ligands per metal ion, the [Si(OH,] in the abscissa term would have to be raised to the z ta power in order to obtain normalization of the data.) Inconsistent results were obtained when ferric solutions more concentrated than 5 x 10-5 M were used. This is probably attributable to the formation of polynuclear complexes. Evaluation of n and Q: from the data given in Fig. 1 is made by rearranging Equation (l) in the logarithmic form log [[FeH.SiO.(--:)+]/[Fe*+][Si(OH).]] = log Qt - (4 - n) log[H +]
(8)
and by plotting the left side of this Equation vs. --log[H+]. The ordinate values in Fig. 1 (inset) were obtained by dividing the quantity [1 + Qa/[H+]] for each series of solutions by the slope of the corresponding trace shown in Fig. I. A value of Q , = 2.89 x 10-s (/z = 0.1, 25°C) (s) has been used in the calculations. From the logarithmic plot values of n = 3 and Q = 0.57 can be derived. Thus, the equilibrium relationships for the complex formation between Fe(III) and orthosilicic acid under the experimental conditions studied ([H +] < 10-s's and [Fe(III)] < 10-4"s) can be represented by Fe s+ + Si(OH), = FeSiO(OH), a+ + H+; Qx = 0.57 (25°C, # = 0.1) t,) T. W. Nv.wa'oN and M. G. ARCAND, J. Amer. Chem. Soc. 75¢ 2449 (1953).
(9)
Notes
239 <3
.ooo ~",.~ /o
350(1
~
3OOO
2,5
[Ferl 2500
I
1
\
iO-IJlO ;5000
3.0
-leg {H+1
I
9-
3
\
4,
5
.A-~
6
7
8
9
,
[Si(OH)4][FeT] {
x I0 "5
Fig. 1
TM
Using 10-.'s as the first concentration acidity constant of Si(OH), the following value of the concentration stability constant for the silicato-ferri complex is obtained: K = Qx/Ksl~om, = [FeSiO(OH)=S+]/[Fe*+][SiO(OH)s-] -~ 1"8 × 10' (25°C,/4 = 0"1)
(10)
It is interesting to compare such complex formation in homogeneous solution with ferric ion and general cation exchange reactions that are ob~rved with dlica gel; ct,loj this ion exchange has been interpreted as surface complex formation of the metal ion with the weakly acidic silanol (--SiOH) groups of the silica gel surface. Natural waters contain as much as 10-s M dis.qolved silica. The results reported here, confirm that iron(IIl) and silicates can interact chemically under certain conditions; such interactior~ have been pmtulated in connection with observations on limnological transformations of iron and silicaJ TM In evaluating the effect of silicates on coagulation by ferric iron ~lt~specific chemical interactions have to be considered in addition to those of the electric double layer. Acknowledgement--This work was supported in part by Public Health Service Grant WP 00043. Division of F~glneering and Applied Physics Harvard University Cambridge 38, Mass.
W . J . Wnleg, Jlt.* W. STUMM
* Present address: Department of Civil Engineering, University of Michigan, Ann Arbor, Michigan. c,~ j. STAWnm and R. W. M^Aa'U~,N,J. Coil. Sci. 18, 132 (1963). ,o~ D. L. DUOOER, J. H. ST~WrON, B. N. I u Y , B. L. McCom~eLL, W. W. Cuunmt,~3s and R. W. MAA'rMAN,J. phys. Chem. 68, 757 (1964). cxx~j. j. MOROANand W. S ~ , in Proe. 2nd Int. Conf. on Water Pollution Control, Tokyo 1964, Pergamon, London. ~ls~j. R. BAYLIS,J. Amer. Water Works Ass. 29, 1355 (1937).