- Email: [email protected]
Solvations of redox ions in glass-forming silicate melts
I O U R N A L OF
ELSEVIER
Journal of Non-Crystalline Solids 181 (1995) 225-230
Solvations of redox ions in glass-forming silicate melts Henry D. Schreiber *, M. Todd Coolbaugh Department of Chemistry, Virginia Military Institute, Lexington, VA 24450, USA Received 5 August 1994
Abstract
The electromotive force series of redox couples have been compared in a variety of solvents including water, methanol, acetonitrile and glass-forming silicate melts. In general, the ordering of reduction potentials of redox couples in aqueous systems is similar to the trend established by the electronegativities of the elements. This ordering of reduction potentials is similar in other non-aqueous solvents, including the silicate melts - indicative of the ability of a species to attract an electron which is the major determinant of its potential to be reduced, regardless of the presence of a solvent. Unique solvations of specific redox ions (Ag +, Cu +, EU2+, Cr 2+, U 5+, and Se 2-) in glass-forming silicate melts can be implied from the deviations in the general trends and can be used to explain the stability of particular redox species.
I. Introduction
Many elements that are introduced into a glass batch can exist in more than one redox state in the melt. For example, iron is typically present as an equilibrium mixture of Fe 2÷ and Fe 3+ ions in a glass-forming melt. The relative concentrations of redox ions are fixed by melt composition, temperature, atmosphere (oxygen fugacity) and the addition of oxidizing or reducing agents [1]. The processing of the glass-forming melt is affected by the multivalent elements present and their redox states. The control of fining, the removal of bubbles from the system, involves an understanding of the oxidation-reduction (redox) reactions that oc-
Presented at the 12th University Conference on Glass Science, Alfred University, Alfred, NY, USA, 25-29 July 1993. * Corresponding author. Tel: + 1-703 464 7244. Telefax: + 1703 464 7261. E-mail: fchschreiber%chemistry%[email protected].
cur within the melt [2]. Likewise, the properties of the glass product such as color and durability can be related to the nature and the amount of the redox ions present in the system. A knowledge of the redox chemistry of the glass-forming melt is thus necessary to produce glass with target characteristics. Redox reactions are chemical reactions that involve the exchange of electrons from one component to another. Although fundamentally a very simple reaction in concept, the mechanics of electron transfer can be quite complex in the presence of a solvent such as a silicate melt. The silicate solvent can be an active participant in establishing its redox state through solvolysis reactions, whereby oxygen can be incorporated into or released from the solvent network by electron exchange reactions with multivalent components of the melt. This is illustrated in Eq. (1) for the case of iron redox equilibria: 4Fe 3+ (melt) + 2 0 2 - (melt)
0022-3093/95/$09.50 © 1995 Elsevier Science B.V. All rights reserved SSD1 0 0 2 2 - 3 0 9 3 ( 9 4 ) 0 0 5 1 6 - 8
'~" 4Fe 2+ (melt) + O~(gas),
(1)
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995) 225-230
226
where the oxide ion represents the polymerization of the silicate network [3]. Depending on the assumed or determined solvation of the iron redox species, this solvolysis reaction can also be expressed as shown in Eq. (2) [4]:
~2
.,,.f
i' 0
4FeO 2 (melt) 4Fe 2+ (melt) + Oa(gas ) + 6 0 2 - (melt)
(2)
In addition, in concert with internal mutual interactions of two redox components dissolved within the system, the solvent molecules must also reorient themselves when an electron is transferred from one site to another. The example of adding cerium to the melt to adjust the iron redox state is shown in Eq.
-4 0.5
I
1.5
2
2,5
3
ELECTRONEGATIVITY (PAULING SCALE)
Fig. 1. The standard reduction potential in aqueous solution for the most common ion of metallic elements versus the Pauling scale electronegativity value of that metal.
(3): Fe 2÷ (melt) + Ce 4+ (melt) Fe 3+ (melt) + Ce 3+ (melt)
(3)
The solvation of the iron redox ion, or the coordination of solvent molecules around the ferrous and ferric ions, changes in the course of this reaction [5].
2. Constraints on theory
The concept of standard reduction potentials has been central in understanding and organizing redox reactions in aqueous solutions [6]. The reduction potentials of redox couples are determined under standard conditions (25°C, 1 atm) with respect to the hydrogen ion-hydrogen couple, for which the reduction of H ÷ ions to molecular H 2 is assigned a value of zero. For example, the half reaction for the Ag ÷ Ag o couple, Ag + (aq) + e - ~ Ag(s)
E ° = + 0.80 V,
(4)
indicates that the silver ion is reduced preferentially to the hydrogen ion in aqueous solution. In addition, the oxidation potential of silver metal in aqueous solution is simply - 0 . 8 0 V, the opposite sign from the reduction of Ag ÷ because of the reversibility of the electron transfer. Values of individual redox potentials are sensitive to changes in temperature, concentration, pH and complexing agents, but models are available to determine the values under nonstandard conditions. The series of reduction potentials of redox couples in water are effectively used to predict reaction
spontaneity. If a net positive potential results upon addition of the reduction potential and the oxidation potential of the two reacting components, the redox reaction would be expected to occur in this solvent. Thus, the power of organizing redox chemistry within the framework of reduction potentials is the capability to determine which reactions do or do not occur or, alternatively, what oxidizing or reducing agent must be added to result in the desired products. Fig. 1 compares the standard reduction potentials of metal ions in water to the elements' electronegativities, a measure of the ability of the metal to attract an electron. The correlation of these two concepts implies that the reduction of a species is primarily controlled by its inherent ability to attract an electron, independent of the presence of the water as solvent [6]. However, this is only a general trend, with the scatter in experimental points in Fig. 1 explained by slightly differing solvations of the various ions within the aqueous environment. The electromotive force series, or listing of standard reduction potentials, in the aqueous environment is compared with that in methanol (CH3OH) as a solvent in Fig. 2 and with that with acetonitrile (CH3CN) as a solvent in Fig. 3 That the series remains relatively solvent-independent is indicative, once again, that the inherent reducibility of a species is the overriding factor in controlling its reduction potential. The values of reduction potentials simply shift en masse from one solvent to another without changing the relative positions of the redox couples within the series. However, exceptions to the general trend are valuable in identifying different solvations
H.D. Schreiber, M.T. Coolbaugh /Journal of Non-Crystalline Solids 181 (1995) 225-230 0
-2
~l
-3 -2 -I 0 REDUCTION POTENTIAL (WATER)
Fig. 2. The standard reduction potentials of various redox couples in methanol versus the standard reduction potentials in water.
of ions in the solvents. For example, one anomalous point in the water-methanol correlation (Fig. 2) is the reduction potential of the A13+-A1 ° couple. That the aluminum ion is much easier to reduce in the methanol environment demonstrates the enhanced stability of the hexahydrated aluminum ion, AI(H 20)63+, in water. Solvent molecules of methanol, being larger than water, cannot coordinate six oxygens around the aluminum ion to stabilize it. For the water-acetonitrile correlation (Fig. 3), the Cu 2+Cu ÷ couple falls somewhat off the general correlation with the reduction being more favorable in the acetonitrile solvent. Similar behavior is also noted in ammonia as a solvent. This is readily explained by the Cu + ion being very effectively solvated by the nitrogen atoms of acetonitrile, consequently destabilizing the Cu 2÷ species in a relative sense in this solvent [7]. The glass-forming silicate melt can be considered to be simply another non-aqueous solvent, such as
2
i° -I
Z-2
-4
L
-4
-3 -2 -1 0 REDUCTION POTENTIAL (WATER)
1
2
Fig. 3. T h e s t a n d a r d r e d u c t i o n potentials o f v a r i o u s r e d o x c o u p l e s in acetonitrile v e r s u s the s t a n d a r d r e d u c t i o n potentials in w a t e r .
227
methanol or acetonitrile, which is a medium for electron exchange reactions. On the one hand, the electromotive force series in a silicate melt would be expected to be similar to that in other solvents because the reduction potentials are primarily controlled by the inherent abilities to attract electrons; but, on the other hand, individual redox couples could be quite sensitive to unique solvations in the melt. Both glass-forming silicate melts and water as solvents are expected to be similar since both are highly structured liquids with their oxygen atoms coordinating to the redox ions. A series of reduction potentials of multivalent elements in a glass-forming melt would be useful to predict the spontaneity of redox reactions in this solvent [8,9]. This electromotive force series could then be used to determine the type and amount of reducing or oxidizing agent(s) necessary to provide the desired final redox state of the melt or glass. Such a series has been determined for a reference borosilicate system [10,11], which also allows comparisons of this series with those in other solvent systems. This approach can aid in understanding the role of the silicate network in solvating redox ions and, thus, in regulating redox reactions within its environment.
3. Results Table 1 provides a comprehensive scale of relative reduction potentials of redox couples in a reference alkali borosilicate glass melt. Reference conditions, which are not standard state conditions, are defined as 1 wt% of the redox component dissolved in a melt of composition SRL-131 at 1150°C and 1 atm. The composition of SRL-131 is that of an alkali borosilicate, one that has been used as a model glass frit not only for nuclear waste immobilization but also for commercial glasses [11]. The reduction potentials for this series were obtained from an experimental analysis of the dependence of the equilibrium redox ratio with the imposed oxygen fugacity [10,11]. The electromotive force series in the reference melt SRL-131 is compared with the series of standard reduction potentials in aqueous solution in Figs. 4 and 5. The good correlation between the two series is indicative once again that the inherent reducibility
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-CrystaUine Solids 181 (1995) 225-230
228
Table 1 A listing of the electromotive force series o f redox couples in the melt S R L - 1 3 1 (composition: 57.9 w t % SiO 2, 1.0% TiO 2, 0.5% ZrO2, 14.7% B 2 0 3 , 0.5% L a x O 3, 2.0% M g O , 17.7% N a 2 0 , and 5.7% L i 2 0 ) at 1150°C and the c o r r e s p o n d i n g standard aqueous reduction potentials R e d o x couple
E' (melt)
A u 3÷ - A u ° p d 2+ - P d ° Ni 3+ - N i 2 + Rh 3+ - R h ° Co 3+ - C o 2+ M n 3+ - M n 2 ÷ Ag ÷ -Ag o
+ 3.0 + 2.3 + 1.7 + 1.4 + 1.4 + 0.8 + 0.5 + 0.2 + 0.2 --0.1 - 0.3 - 0.3 -0.8 - 1.5 - 1.7 - 1.7 - 1.9 - 2.2
+ 1.42 + 0.92 + 1.75 + 0.76 + 1.83 + 1.51 + 0.80 + 0.88 + 0.88 + 1.44 + 1.22 + 0.68 +0.16 +0.06 +0.77 + 0.58 + 1.00 + 0.61
- 3.3 -3.4 - 3.8 -4.0 - 4.3 - 5.0 - 5.3 - 5.5 - 5.7 -6.0 - 6.3 - 15.8 - 19.2
+ 0.52 -0.41 + 0.48 +0.34 - 0.35 + 0.03 - 0.24 + O. 14 - 0.92 -0.28 - 0.41
S e 6+ - S e 4+
Se 4+ - S e ° Ce 4+ - C e 3+ Cr 6 + - C r 3 + Sb 5 + - S b 3 + Cu2+-Cu + U 6+ - U s +
Fe 3+ - F e E+ A s 5+ - A s 3 + V 5+ - V 4+ U 5+ - U 4 + Cu + - C u ° C r 3 + - C r 2+ M o 6+ - M o 5 + V 4+ - V 3+ Eu 3 + - E u 2 + Ti 4+ -Ti3 + Ni 2÷ - N i ° Sn 4 + - S n 2 + Se ° - S e 2 Co2+-Co ° Fe 2 + - F e ° Mo 5+ - M o O S 6+ - S 2 -
E ° (water)
U6+-U
5+,
EU3+-Eu
S 0 0.5 1 1.5 REDUCTION POTENTIAL (WATER)
versus the standard reduction potentials in water.
A4 o-i
I O
-8
•
I
-1
.'.5
o
015
REDUCTION POTENTIAL 0~ATER)
,
,.5
Fig. 5. The relative reduction potentials of various multiple-electron redox couples in the melt S R L - 1 3 1 u n d e r reference conditions versus the standard reduction potentials in water.
the reference system SRL-131. This system has a composition within the forsterite-anorthite-silica (FAS) phase diagram and, as such, is a magnesiumcalcium-aluminosilicate. This correlation is quite
the silicate melt than in the aqueous system, namely +,
-2 ~-3 •
Fig. 4. The relative reduction potentials o f various one-electron redox couples in the melt S R L - 1 3 1 u n d e r reference conditions
of a species controls its reduction potential, whether in water or in the glass-forming melt. However, there are some redox couples which are easier to reduce in Cu2+-Cu
0
-6 -0.5
The relative reduction potentials ( E ' ) in S R L - 1 3 1 w e r e determ i n e d f r o m the equation l o g ( X ) = ( n / 4 ) ( - l o g fo2)+E', where X is the ratio o f the concentrations of the element in the reduced state to the oxidized state, n is the n u m b e r o f electrons transferred in the redox couple, and fo2 is the imposed o x y g e n fugacity.
Ag+-Ag °,
2
2+,
Cr3+-Cr 2+, and Se°-Se 2-. The redox couple Cr 6+Cr 3+ is marginally harder to reduce in the silicate melt than in the aqueous solvent. Fig. 6 compares the scale of reduction potentials in another glass-forming melt at 1500°C with that in
~
2 1
-2
1_, -6
-5 -4 -3 -2 -1 REDUCTION POTENTIAL (SRL-131)
Fig. 6. The relative reduction potentials of various redox couples in the melt F A S (54.1 w t % SiO2, 10.8% A 1 2 0 3 , 5.7% C a O , and 2 9 . 4 % M g O ) at 1500°C versus the relative reduction potentials in the melt S R L - 1 3 1 u n d e r reference conditions.
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995) 225-230
good, as also is the one with sodium disilicate as a solvent at 1080°C [10]. This study determined the relative reduction potentials in the reference glass-forming melt through an indirect procedure, by experimentally measuring the equilibrium redox ratios of the individual elements as a function of the imposed oxygen fugacity. It is also possible to obtain such reduction potentials directly by electrochemical techniques [12], although a comprehensive series in a particular silicate solvent is not yet available by this procedure. However, the series compiled in Table 1 for the reference system SRL-131 has been shown to be in agreement with that determined electrochemically in another silicate melt system [13].
4. Discussion The correspondence between the electromotive force series in the reference melt versus that in water indicates that the solvations of most redox ions are analogous in both solvents. This provides credence to the use of, for example, comparisons with absorption spectra in water to ascertain the coordination sites of redox ions in glass melts [14]. In both solvents, the solvation of the ion is to the oxygen atoms within the structured solvent network. However, some unique solvations or stabilities of redox ions in glass melts have been identified. When a particular redox couple is more easily reduced in the glass melt than in the aqueous solvent, either the solvation of the oxidized ion is destabilized in the melt or the solvation of the reduced ion is more stabilized in the melt. Ag ÷ ions are more stable in aqueous solution than in glass melts, owing to the reduction of this ion at higher temperatures and the known incompatibility of this ion in silicate networks. U 6÷ ions, as the uranyl UO 2+ species, are easier to reduce to U 5+ ions, as the UO~- species, in the silicate network than in aqueous solution. The silicate network stabilizes this U 5÷ entity [15], whereas in water the pentavalent uranium ion readily undergoes disproportionation reactions [7]. Divalent Eu and Cr are also stabilized in the glass melt because of their ready solvation in the melt versus aqueous systems. In aqueous solutions both Eu z÷ and Cr 2÷ have similar redox chemistries [7]. Univa-
229
lent copper is stabilized in the glass melt, much like in acetonitrile, as compared with aqueous solutions. The relative stabilities of Cu 2+ and Cu ÷ even in aqueous solutions are known to depend very strongly on the nature of anions or ligands present and, as such, vary considerably with the solvent [7]. In addition, the univalent form of copper disproportionates somewhat in water. The selenide ion, probably due to its substitution for the oxide ion in the network, is stabilized in the melt environment with respect to the elemental selenium form. Finally, the trivalent form of chromium is less stable in the glass melt than in aqueous solutions, as illustrated by both the Cr 6+ (chromate)-Cr 3÷ and Cr3+-Cr2+ redox couples. Since the trivalent state of chromium has a great desire for octahedral coordination, water more effectively solvates this ion than the silicate network. Interestingly, the deviations from the correspondence of the electromotive force series in the solvents (Figs. 4 and 5) only identify major changes in solvation from one system to another. For example, Co 2+ in water possesses octahedral coordination and introduces a pink coloration to the solution, while in the reference melt, the ion is tetrahedrally coordinated with a blue coloration. However, the reduction of Co 2+ ions to Co metal does not seem to be anomalous in the melt as opposed to water. The energy differences between these two coordinations of Co 2+ are very small and evidently are not manifested in the correlation of Fig. 5. Use of the melt-water correlations of the electromotive force series in the two solvents has been applied previously to determine the redox chemistry of species which are difficult to measure experimentally in silicate melts. For example, the redox chemistry and solubility of precious metals in melts [16], as well as the chemistry of plutonium and other radioactive species in melts [10], have been extrapolated from such relations. The choice of redox agents to obtain the desired redox state in the final glass product can accordingly be determined from the electromotive force series summarized in Table 1.
5. Conclusions The principal factor affecting the reduction potential of a redox couple is the inherent ability of the
230
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995)225-230
species to attract an electron; however, solvations of the ions also play a role in the reduction potential since selective solvation can stabilize or destabilize a particular redox ion of that couple. Silicate melts solvate certain ions such as Cu ÷ and U 5÷ more effectively than does water but, on the other hand, they solvate ions such as Ag ÷ and Cr 3+ less effectively. Identification of unique solvations of these ions will aid in the understanding of internal mutual interactions which involve such ions and thus changes in solvation spheres in the course of the redox reactions. Support for this research has been furnished by the NSF Industry-University Center for Glass Chemistry (Alfred University, New York) as well as the Camille and Henry Dreyfus Scholar/Fellow program.
References [1] H.D. Schreiber, J. Non-Cryst. Solids 84 (1986) 129. [2] H.D. Schreiber, P.G. Leonhard, R.G. Nofsinger, M.W. Henning, C.W. Schreiber and S.J. Kozak, in: Advances in the
[3] [4] [5] [6] [7]
[8] [9] [10] [11]
[12] [13] [14] [15] [16]
Fusion of Glass, ed. D.F. Bickford et al. (American Ceramic Society, Westerville, OH, 1988) p. 29.1. K.H. Karlsson, L. FrBberg and B. Skrifvars, Glastech. Ber. 63 (1990) 37. H.D. Schreiber and G.B. Balazs, J. Non-Cryst. Solids 71 (1985) 59. B.O. Mysen, in: Physical Chemistry of Magmas, ed. L.L. Perchuk and I. Kushiro (Springer, New York, 1991) p. 41. G. Wulfsberg, Principles of Descriptive Inorganic Chemistry (Brooks/Cole, Monterey, CA, 1987) p. 137. F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 4th Ed. (Wiley, New York, 1980) pp. 801, 1002 and 1034. H.D. Schreiber, R.C. Merkel, V.L. Schreiber and G.B. Balazs, J. Geophys. Res. 92 (1987) 9283. H.D. Schreiber and A.L. Hockman, J. Am. Ceram. Soc. 70 (1987) 591. H.D. Schreiber, J. Geophys. Res. 70 (1987) 9225. H.D. Schreiber, G.B. Balazs, B.E. Carpenter, J.E. Kirkley, L.M. Minnix and P.L. Jamison, J. Am. Ceram. Soc. 67 (1984) C106. C. Riissel and E. Freude, Phys. Chem. Glasses 30 (1989) 62. R.O. Colson, L.A. Haskin and D. Crane, Geochim. Cosmochim. Acta 54 (1990) 3353. C.R. Bamford, Phys. Chem. Glasses 3 (1962) 189. H.D. Schreiber and G.B. Balazs, Phys. Chem. Glasses 23 (1982) 147. H.D. Schreiber, T.R. Harville and G.N. Damron, J. Am. Ceram. Soc. 73 (1990) 1435.
ELSEVIER
Journal of Non-Crystalline Solids 181 (1995) 225-230
Solvations of redox ions in glass-forming silicate melts Henry D. Schreiber *, M. Todd Coolbaugh Department of Chemistry, Virginia Military Institute, Lexington, VA 24450, USA Received 5 August 1994
Abstract
The electromotive force series of redox couples have been compared in a variety of solvents including water, methanol, acetonitrile and glass-forming silicate melts. In general, the ordering of reduction potentials of redox couples in aqueous systems is similar to the trend established by the electronegativities of the elements. This ordering of reduction potentials is similar in other non-aqueous solvents, including the silicate melts - indicative of the ability of a species to attract an electron which is the major determinant of its potential to be reduced, regardless of the presence of a solvent. Unique solvations of specific redox ions (Ag +, Cu +, EU2+, Cr 2+, U 5+, and Se 2-) in glass-forming silicate melts can be implied from the deviations in the general trends and can be used to explain the stability of particular redox species.
I. Introduction
Many elements that are introduced into a glass batch can exist in more than one redox state in the melt. For example, iron is typically present as an equilibrium mixture of Fe 2÷ and Fe 3+ ions in a glass-forming melt. The relative concentrations of redox ions are fixed by melt composition, temperature, atmosphere (oxygen fugacity) and the addition of oxidizing or reducing agents [1]. The processing of the glass-forming melt is affected by the multivalent elements present and their redox states. The control of fining, the removal of bubbles from the system, involves an understanding of the oxidation-reduction (redox) reactions that oc-
Presented at the 12th University Conference on Glass Science, Alfred University, Alfred, NY, USA, 25-29 July 1993. * Corresponding author. Tel: + 1-703 464 7244. Telefax: + 1703 464 7261. E-mail: fchschreiber%chemistry%[email protected].
cur within the melt [2]. Likewise, the properties of the glass product such as color and durability can be related to the nature and the amount of the redox ions present in the system. A knowledge of the redox chemistry of the glass-forming melt is thus necessary to produce glass with target characteristics. Redox reactions are chemical reactions that involve the exchange of electrons from one component to another. Although fundamentally a very simple reaction in concept, the mechanics of electron transfer can be quite complex in the presence of a solvent such as a silicate melt. The silicate solvent can be an active participant in establishing its redox state through solvolysis reactions, whereby oxygen can be incorporated into or released from the solvent network by electron exchange reactions with multivalent components of the melt. This is illustrated in Eq. (1) for the case of iron redox equilibria: 4Fe 3+ (melt) + 2 0 2 - (melt)
0022-3093/95/$09.50 © 1995 Elsevier Science B.V. All rights reserved SSD1 0 0 2 2 - 3 0 9 3 ( 9 4 ) 0 0 5 1 6 - 8
'~" 4Fe 2+ (melt) + O~(gas),
(1)
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995) 225-230
226
where the oxide ion represents the polymerization of the silicate network [3]. Depending on the assumed or determined solvation of the iron redox species, this solvolysis reaction can also be expressed as shown in Eq. (2) [4]:
~2
.,,.f
i' 0
4FeO 2 (melt) 4Fe 2+ (melt) + Oa(gas ) + 6 0 2 - (melt)
(2)
In addition, in concert with internal mutual interactions of two redox components dissolved within the system, the solvent molecules must also reorient themselves when an electron is transferred from one site to another. The example of adding cerium to the melt to adjust the iron redox state is shown in Eq.
-4 0.5
I
1.5
2
2,5
3
ELECTRONEGATIVITY (PAULING SCALE)
Fig. 1. The standard reduction potential in aqueous solution for the most common ion of metallic elements versus the Pauling scale electronegativity value of that metal.
(3): Fe 2÷ (melt) + Ce 4+ (melt) Fe 3+ (melt) + Ce 3+ (melt)
(3)
The solvation of the iron redox ion, or the coordination of solvent molecules around the ferrous and ferric ions, changes in the course of this reaction [5].
2. Constraints on theory
The concept of standard reduction potentials has been central in understanding and organizing redox reactions in aqueous solutions [6]. The reduction potentials of redox couples are determined under standard conditions (25°C, 1 atm) with respect to the hydrogen ion-hydrogen couple, for which the reduction of H ÷ ions to molecular H 2 is assigned a value of zero. For example, the half reaction for the Ag ÷ Ag o couple, Ag + (aq) + e - ~ Ag(s)
E ° = + 0.80 V,
(4)
indicates that the silver ion is reduced preferentially to the hydrogen ion in aqueous solution. In addition, the oxidation potential of silver metal in aqueous solution is simply - 0 . 8 0 V, the opposite sign from the reduction of Ag ÷ because of the reversibility of the electron transfer. Values of individual redox potentials are sensitive to changes in temperature, concentration, pH and complexing agents, but models are available to determine the values under nonstandard conditions. The series of reduction potentials of redox couples in water are effectively used to predict reaction
spontaneity. If a net positive potential results upon addition of the reduction potential and the oxidation potential of the two reacting components, the redox reaction would be expected to occur in this solvent. Thus, the power of organizing redox chemistry within the framework of reduction potentials is the capability to determine which reactions do or do not occur or, alternatively, what oxidizing or reducing agent must be added to result in the desired products. Fig. 1 compares the standard reduction potentials of metal ions in water to the elements' electronegativities, a measure of the ability of the metal to attract an electron. The correlation of these two concepts implies that the reduction of a species is primarily controlled by its inherent ability to attract an electron, independent of the presence of the water as solvent [6]. However, this is only a general trend, with the scatter in experimental points in Fig. 1 explained by slightly differing solvations of the various ions within the aqueous environment. The electromotive force series, or listing of standard reduction potentials, in the aqueous environment is compared with that in methanol (CH3OH) as a solvent in Fig. 2 and with that with acetonitrile (CH3CN) as a solvent in Fig. 3 That the series remains relatively solvent-independent is indicative, once again, that the inherent reducibility of a species is the overriding factor in controlling its reduction potential. The values of reduction potentials simply shift en masse from one solvent to another without changing the relative positions of the redox couples within the series. However, exceptions to the general trend are valuable in identifying different solvations
H.D. Schreiber, M.T. Coolbaugh /Journal of Non-Crystalline Solids 181 (1995) 225-230 0
-2
~l
-3 -2 -I 0 REDUCTION POTENTIAL (WATER)
Fig. 2. The standard reduction potentials of various redox couples in methanol versus the standard reduction potentials in water.
of ions in the solvents. For example, one anomalous point in the water-methanol correlation (Fig. 2) is the reduction potential of the A13+-A1 ° couple. That the aluminum ion is much easier to reduce in the methanol environment demonstrates the enhanced stability of the hexahydrated aluminum ion, AI(H 20)63+, in water. Solvent molecules of methanol, being larger than water, cannot coordinate six oxygens around the aluminum ion to stabilize it. For the water-acetonitrile correlation (Fig. 3), the Cu 2+Cu ÷ couple falls somewhat off the general correlation with the reduction being more favorable in the acetonitrile solvent. Similar behavior is also noted in ammonia as a solvent. This is readily explained by the Cu + ion being very effectively solvated by the nitrogen atoms of acetonitrile, consequently destabilizing the Cu 2÷ species in a relative sense in this solvent [7]. The glass-forming silicate melt can be considered to be simply another non-aqueous solvent, such as
2
i° -I
Z-2
-4
L
-4
-3 -2 -1 0 REDUCTION POTENTIAL (WATER)
1
2
Fig. 3. T h e s t a n d a r d r e d u c t i o n potentials o f v a r i o u s r e d o x c o u p l e s in acetonitrile v e r s u s the s t a n d a r d r e d u c t i o n potentials in w a t e r .
227
methanol or acetonitrile, which is a medium for electron exchange reactions. On the one hand, the electromotive force series in a silicate melt would be expected to be similar to that in other solvents because the reduction potentials are primarily controlled by the inherent abilities to attract electrons; but, on the other hand, individual redox couples could be quite sensitive to unique solvations in the melt. Both glass-forming silicate melts and water as solvents are expected to be similar since both are highly structured liquids with their oxygen atoms coordinating to the redox ions. A series of reduction potentials of multivalent elements in a glass-forming melt would be useful to predict the spontaneity of redox reactions in this solvent [8,9]. This electromotive force series could then be used to determine the type and amount of reducing or oxidizing agent(s) necessary to provide the desired final redox state of the melt or glass. Such a series has been determined for a reference borosilicate system [10,11], which also allows comparisons of this series with those in other solvent systems. This approach can aid in understanding the role of the silicate network in solvating redox ions and, thus, in regulating redox reactions within its environment.
3. Results Table 1 provides a comprehensive scale of relative reduction potentials of redox couples in a reference alkali borosilicate glass melt. Reference conditions, which are not standard state conditions, are defined as 1 wt% of the redox component dissolved in a melt of composition SRL-131 at 1150°C and 1 atm. The composition of SRL-131 is that of an alkali borosilicate, one that has been used as a model glass frit not only for nuclear waste immobilization but also for commercial glasses [11]. The reduction potentials for this series were obtained from an experimental analysis of the dependence of the equilibrium redox ratio with the imposed oxygen fugacity [10,11]. The electromotive force series in the reference melt SRL-131 is compared with the series of standard reduction potentials in aqueous solution in Figs. 4 and 5. The good correlation between the two series is indicative once again that the inherent reducibility
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-CrystaUine Solids 181 (1995) 225-230
228
Table 1 A listing of the electromotive force series o f redox couples in the melt S R L - 1 3 1 (composition: 57.9 w t % SiO 2, 1.0% TiO 2, 0.5% ZrO2, 14.7% B 2 0 3 , 0.5% L a x O 3, 2.0% M g O , 17.7% N a 2 0 , and 5.7% L i 2 0 ) at 1150°C and the c o r r e s p o n d i n g standard aqueous reduction potentials R e d o x couple
E' (melt)
A u 3÷ - A u ° p d 2+ - P d ° Ni 3+ - N i 2 + Rh 3+ - R h ° Co 3+ - C o 2+ M n 3+ - M n 2 ÷ Ag ÷ -Ag o
+ 3.0 + 2.3 + 1.7 + 1.4 + 1.4 + 0.8 + 0.5 + 0.2 + 0.2 --0.1 - 0.3 - 0.3 -0.8 - 1.5 - 1.7 - 1.7 - 1.9 - 2.2
+ 1.42 + 0.92 + 1.75 + 0.76 + 1.83 + 1.51 + 0.80 + 0.88 + 0.88 + 1.44 + 1.22 + 0.68 +0.16 +0.06 +0.77 + 0.58 + 1.00 + 0.61
- 3.3 -3.4 - 3.8 -4.0 - 4.3 - 5.0 - 5.3 - 5.5 - 5.7 -6.0 - 6.3 - 15.8 - 19.2
+ 0.52 -0.41 + 0.48 +0.34 - 0.35 + 0.03 - 0.24 + O. 14 - 0.92 -0.28 - 0.41
S e 6+ - S e 4+
Se 4+ - S e ° Ce 4+ - C e 3+ Cr 6 + - C r 3 + Sb 5 + - S b 3 + Cu2+-Cu + U 6+ - U s +
Fe 3+ - F e E+ A s 5+ - A s 3 + V 5+ - V 4+ U 5+ - U 4 + Cu + - C u ° C r 3 + - C r 2+ M o 6+ - M o 5 + V 4+ - V 3+ Eu 3 + - E u 2 + Ti 4+ -Ti3 + Ni 2÷ - N i ° Sn 4 + - S n 2 + Se ° - S e 2 Co2+-Co ° Fe 2 + - F e ° Mo 5+ - M o O S 6+ - S 2 -
E ° (water)
U6+-U
5+,
EU3+-Eu
S 0 0.5 1 1.5 REDUCTION POTENTIAL (WATER)
versus the standard reduction potentials in water.
A4 o-i
I O
-8
•
I
-1
.'.5
o
015
REDUCTION POTENTIAL 0~ATER)
,
,.5
Fig. 5. The relative reduction potentials of various multiple-electron redox couples in the melt S R L - 1 3 1 u n d e r reference conditions versus the standard reduction potentials in water.
the reference system SRL-131. This system has a composition within the forsterite-anorthite-silica (FAS) phase diagram and, as such, is a magnesiumcalcium-aluminosilicate. This correlation is quite
the silicate melt than in the aqueous system, namely +,
-2 ~-3 •
Fig. 4. The relative reduction potentials o f various one-electron redox couples in the melt S R L - 1 3 1 u n d e r reference conditions
of a species controls its reduction potential, whether in water or in the glass-forming melt. However, there are some redox couples which are easier to reduce in Cu2+-Cu
0
-6 -0.5
The relative reduction potentials ( E ' ) in S R L - 1 3 1 w e r e determ i n e d f r o m the equation l o g ( X ) = ( n / 4 ) ( - l o g fo2)+E', where X is the ratio o f the concentrations of the element in the reduced state to the oxidized state, n is the n u m b e r o f electrons transferred in the redox couple, and fo2 is the imposed o x y g e n fugacity.
Ag+-Ag °,
2
2+,
Cr3+-Cr 2+, and Se°-Se 2-. The redox couple Cr 6+Cr 3+ is marginally harder to reduce in the silicate melt than in the aqueous solvent. Fig. 6 compares the scale of reduction potentials in another glass-forming melt at 1500°C with that in
~
2 1
-2
1_, -6
-5 -4 -3 -2 -1 REDUCTION POTENTIAL (SRL-131)
Fig. 6. The relative reduction potentials of various redox couples in the melt F A S (54.1 w t % SiO2, 10.8% A 1 2 0 3 , 5.7% C a O , and 2 9 . 4 % M g O ) at 1500°C versus the relative reduction potentials in the melt S R L - 1 3 1 u n d e r reference conditions.
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995) 225-230
good, as also is the one with sodium disilicate as a solvent at 1080°C [10]. This study determined the relative reduction potentials in the reference glass-forming melt through an indirect procedure, by experimentally measuring the equilibrium redox ratios of the individual elements as a function of the imposed oxygen fugacity. It is also possible to obtain such reduction potentials directly by electrochemical techniques [12], although a comprehensive series in a particular silicate solvent is not yet available by this procedure. However, the series compiled in Table 1 for the reference system SRL-131 has been shown to be in agreement with that determined electrochemically in another silicate melt system [13].
4. Discussion The correspondence between the electromotive force series in the reference melt versus that in water indicates that the solvations of most redox ions are analogous in both solvents. This provides credence to the use of, for example, comparisons with absorption spectra in water to ascertain the coordination sites of redox ions in glass melts [14]. In both solvents, the solvation of the ion is to the oxygen atoms within the structured solvent network. However, some unique solvations or stabilities of redox ions in glass melts have been identified. When a particular redox couple is more easily reduced in the glass melt than in the aqueous solvent, either the solvation of the oxidized ion is destabilized in the melt or the solvation of the reduced ion is more stabilized in the melt. Ag ÷ ions are more stable in aqueous solution than in glass melts, owing to the reduction of this ion at higher temperatures and the known incompatibility of this ion in silicate networks. U 6÷ ions, as the uranyl UO 2+ species, are easier to reduce to U 5+ ions, as the UO~- species, in the silicate network than in aqueous solution. The silicate network stabilizes this U 5÷ entity [15], whereas in water the pentavalent uranium ion readily undergoes disproportionation reactions [7]. Divalent Eu and Cr are also stabilized in the glass melt because of their ready solvation in the melt versus aqueous systems. In aqueous solutions both Eu z÷ and Cr 2÷ have similar redox chemistries [7]. Univa-
229
lent copper is stabilized in the glass melt, much like in acetonitrile, as compared with aqueous solutions. The relative stabilities of Cu 2+ and Cu ÷ even in aqueous solutions are known to depend very strongly on the nature of anions or ligands present and, as such, vary considerably with the solvent [7]. In addition, the univalent form of copper disproportionates somewhat in water. The selenide ion, probably due to its substitution for the oxide ion in the network, is stabilized in the melt environment with respect to the elemental selenium form. Finally, the trivalent form of chromium is less stable in the glass melt than in aqueous solutions, as illustrated by both the Cr 6+ (chromate)-Cr 3÷ and Cr3+-Cr2+ redox couples. Since the trivalent state of chromium has a great desire for octahedral coordination, water more effectively solvates this ion than the silicate network. Interestingly, the deviations from the correspondence of the electromotive force series in the solvents (Figs. 4 and 5) only identify major changes in solvation from one system to another. For example, Co 2+ in water possesses octahedral coordination and introduces a pink coloration to the solution, while in the reference melt, the ion is tetrahedrally coordinated with a blue coloration. However, the reduction of Co 2+ ions to Co metal does not seem to be anomalous in the melt as opposed to water. The energy differences between these two coordinations of Co 2+ are very small and evidently are not manifested in the correlation of Fig. 5. Use of the melt-water correlations of the electromotive force series in the two solvents has been applied previously to determine the redox chemistry of species which are difficult to measure experimentally in silicate melts. For example, the redox chemistry and solubility of precious metals in melts [16], as well as the chemistry of plutonium and other radioactive species in melts [10], have been extrapolated from such relations. The choice of redox agents to obtain the desired redox state in the final glass product can accordingly be determined from the electromotive force series summarized in Table 1.
5. Conclusions The principal factor affecting the reduction potential of a redox couple is the inherent ability of the
230
H.D. Schreiber, M.T. Coolbaugh/Journal of Non-Crystalline Solids 181 (1995)225-230
species to attract an electron; however, solvations of the ions also play a role in the reduction potential since selective solvation can stabilize or destabilize a particular redox ion of that couple. Silicate melts solvate certain ions such as Cu ÷ and U 5÷ more effectively than does water but, on the other hand, they solvate ions such as Ag ÷ and Cr 3+ less effectively. Identification of unique solvations of these ions will aid in the understanding of internal mutual interactions which involve such ions and thus changes in solvation spheres in the course of the redox reactions. Support for this research has been furnished by the NSF Industry-University Center for Glass Chemistry (Alfred University, New York) as well as the Camille and Henry Dreyfus Scholar/Fellow program.
References [1] H.D. Schreiber, J. Non-Cryst. Solids 84 (1986) 129. [2] H.D. Schreiber, P.G. Leonhard, R.G. Nofsinger, M.W. Henning, C.W. Schreiber and S.J. Kozak, in: Advances in the
[3] [4] [5] [6] [7]
[8] [9] [10] [11]
[12] [13] [14] [15] [16]
Fusion of Glass, ed. D.F. Bickford et al. (American Ceramic Society, Westerville, OH, 1988) p. 29.1. K.H. Karlsson, L. FrBberg and B. Skrifvars, Glastech. Ber. 63 (1990) 37. H.D. Schreiber and G.B. Balazs, J. Non-Cryst. Solids 71 (1985) 59. B.O. Mysen, in: Physical Chemistry of Magmas, ed. L.L. Perchuk and I. Kushiro (Springer, New York, 1991) p. 41. G. Wulfsberg, Principles of Descriptive Inorganic Chemistry (Brooks/Cole, Monterey, CA, 1987) p. 137. F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 4th Ed. (Wiley, New York, 1980) pp. 801, 1002 and 1034. H.D. Schreiber, R.C. Merkel, V.L. Schreiber and G.B. Balazs, J. Geophys. Res. 92 (1987) 9283. H.D. Schreiber and A.L. Hockman, J. Am. Ceram. Soc. 70 (1987) 591. H.D. Schreiber, J. Geophys. Res. 70 (1987) 9225. H.D. Schreiber, G.B. Balazs, B.E. Carpenter, J.E. Kirkley, L.M. Minnix and P.L. Jamison, J. Am. Ceram. Soc. 67 (1984) C106. C. Riissel and E. Freude, Phys. Chem. Glasses 30 (1989) 62. R.O. Colson, L.A. Haskin and D. Crane, Geochim. Cosmochim. Acta 54 (1990) 3353. C.R. Bamford, Phys. Chem. Glasses 3 (1962) 189. H.D. Schreiber and G.B. Balazs, Phys. Chem. Glasses 23 (1982) 147. H.D. Schreiber, T.R. Harville and G.N. Damron, J. Am. Ceram. Soc. 73 (1990) 1435.