Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts

Geochimica et Cosmochimica Acta, Vol. 65, No. 11, pp. 1809 –1819, 2001 Copyright © 2001 Elsevier Science Ltd Printed in the USA. All rights reserved 0016-7037/01 $20.00 ⫹ .00

Pergamon

PII S0016-7037(00)00561-5

Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts JEAN A. TANGEMAN,* REBECCA LANGE, and LAURA FORMAN Department of Geological Sciences, 2534 C.C. Little Building, University of Michigan, Ann Arbor, MI 48109-1063, USA (Received February 1, 2000; accepted in revised form August 2, 2000)

Abstract—Seventy-five K2O-FeO-Fe2O3-SiO2 liquids were equilibrated in air between 1123°C and 1596°C. The bulk compositions of the quenched glasses range from 35 to 85 wt.% SiO2, 3 to 41 wt.% Fe2O3, 0 to 5 wt.% FeO, and 11 to 34 wt.% K2O. Despite the volatilization of significant amounts of potassium in many of the high-temperature runs, no gradients in K2O, SiO2, or total iron were observed from electron microprobe transects of the 75 quenched glass beads. This result strongly suggests that the rate of K diffusion in the melt is faster than the rate of K volatilization at the gas/melt interface. The individual quenched glass beads were homogenous in their ferric-ferrous ratios as well, which further suggests that iron-redox equilibrium was achieved in these experiments. The ferric-ferrous data were used to derive a symmetric, regular solution model, with a standard error of 0.18 wt.% in the prediction of FeO in peralkaline K2O-FeO-Fe2O3-SiO2 liquids. Compositional trends for redox equilibrium in the potassic silicate liquids contrast with those found in sodic and calcic silicate liquids, as well as magmatic liquids. The effect of increasing K2O is to decrease the ferric-ferrous ratio in peralkaline, iron-rich (⬎3 wt.%), K2O-FeO-Fe2O3-SiO2 liquids. These results are discussed in terms of three compositional effects that influence iron-redox equilibria in silicate melts: 1) charge-balancing of Fe3⫹; 2) coordination change of Fe3⫹; and 3) the activity of O2⫺ in the melt. Copyright © 2001 Elsevier Science Ltd enable an examination of how the ferric-ferrous ratio varies as a systematic function of K2O, Na2O, CaO, and total iron. A second goal of this study is to develop a thermodynamic model to describe ferric-ferrous equilibria in K2O-FeO-Fe2O3SiO2 liquids as a function of temperature and composition in air. Such a model is necessary before viscosity, heat capacity, density, sound speed, or other property measurements on K2OFeO-Fe2O3-SiO2 liquids can be interpreted adequately. The reason resides in the strong temperature and compositional dependence of the ferric-ferrous ratio in silicate melts. This requires that the redox state of any iron-bearing experimental melt be known at each temperature of measurement. Because it is not often possible to directly sample experimental liquids during property measurements at high temperature, the appropriate ferric-ferrous ratios must be calculated by a well-calibrated model.

1. INTRODUCTION

The structural role of ferric and ferrous iron in magmatic liquids strongly influences how the iron-redox ratio varies as a function of composition, temperature, pressure, and oxygen fugacity. Because of the petrological significance of the ferricferrous ratio as an indicator of oxygen fugacity (e.g., Carmichael, 1991), several experimental studies have explored how it varies in magmatic liquids as a function of these intensive variables (e.g., Fudali, 1965; Thornber et al., 1980; Sack et al., 1980; Kilinc et al., 1983; Kress and Carmichael, 1988; 1991). Although an empirical model has been developed that adequately describes the iron-redox equilibria in magmatic liquids (e.g., Kress and Carmichael, 1991), it is clear that there is significant nonideal mixing between the ferric and ferrous oxide components in natural liquids (Kress and Carmichael, 1988). Understanding the structural basis to this nonideal behavior is therefore of considerable interest to petrologists (e.g., Mysen, 1987). One strategy is to explore the iron-redox equilibria in simple three-to-four component liquids as a supplement to studies on natural compositions (e.g., Paul and Douglas, 1965; Dickenson and Hess, 1981; Mysen et al., 1984; Lange and Carmichael, 1989; Kress and Carmichael, 1989). In these simple systems, the effect of individual cations on the iron-redox equilibria can be isolated. Therefore, one of the goals of this study is to explore the ferric-ferrous equilibria in iron-rich (3–39 wt.% FeOT) potassic silicate liquids, and to compare the redox behavior to that in iron-poor (ⱕ1 wt.% FeOT) potassic silicate liquids (Paul and Douglas, 1965), as well as to the redox equilibria in iron-rich sodic and calcic systems (Lange and Carmichael, 1989; Kress and Carmichael, 1989). This will

2. EXPERIMENTAL PROCEDURE 2.1. Sample Synthesis Starting materials were prepared by weighing out appropriate proportions of reagent grade SiO2, Fe2O3, and K2CO3 powders and vigorously mixing them in a large jar. The powdered mixtures were then transferred to a 200-mL Pt crucible and put into a box furnace at an initial starting temperature of ⬃500°C. The temperature was then slowly raised in incremental steps of 100° (after holding for ⬃30 min at each step) during the decarbonation and melting process to a final temperature of ⬃1100°C. The crucible was removed at ⬃1100°C and the silicate melt was quenched to a glass by immediately immersing the bottom portion of the crucible in ambient temperature water. The glass chunks were ground to a fine powder in an agate ball-mill before being re-fused at 1200°C for 2 to 3 h and quenched to a glass again. The grinding and re-fusion steps were repeated to ensure sample homogeneity. Eight starting glass compositions were synthesized and labeled A, B, C, D, E, F, G, and H (Table 1). A ninth sample (L) was synthesized (fused at 1300°C) to coincide with stoichiometric KFe3⫹Si2O6 (iron-leucite), the only deviation from the mineral being the variable ferric-ferrous ratio in the liquid at various temperatures.

* Author to whom correspondence should be addressed (tangeman@ containerless.com). 1809

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J. A. Tangeman, R. Lange, and L. Forman Table 1. Compositional analyses of the starting glasses.

Sample

SiO2 (wt.%)

Fe2O3 (wt.%)

K2O (wt.%)

Total

KFS-A KFS-B KFS-C KFS-Da KFS-E KFS-F KFS-Gb KFS-H Fe-leuciteb

41.96 35.10 51.69 28.44 63.02 43.25 64.99 53.16 50.41

14.07 23.70 17.94 32.69 20.83 18.67 3.05 9.78 31.19

43.76 40.66 31.05 38.87 15.56 37.52 31.15 37.37 18.42

99.79 99.46 100.68 100.00 99.41 99.44 99.19 100.31 100.02

a Nominal composition; sample was too hygroscopic for wet chemical or microprobe analysis. b Compositions determined from electron microprobe analyses.

intercept method (regression through measurements for 5 counting periods, 2 s each). A plot of these count rates vs. time are shown in Figure 1. Over the 10-s interval, the potassium loss or migration in the samples was linear. The number of counts per second vs. time using a 30-kV accelerating potential and a current of 3 nA were performed on all three standards as well (Fig. 1) and confirmed that no elemental migration in the standards occurred under these conditions. Analytical uncertainties (1␴) for microprobe analyses are given in Table 2. A 2- to 6-mg portion of each quenched sample was used for colorimetric determination of ferrous iron as described by Wilson (1960) and Lange and Carmichael (1989). All samples were first examined for the presence of quench crystals or magnetic behavior. The colorimetric technique, which has a 2␴ precision of ⫾0.21 wt.% FeO, was performed at the University of Michigan on all samples, including those that had crystallized. Replicate measurements were made to ensure homogeneity with respect to the ferrous iron content. Ferric iron was determined by difference, where Fe2O3 wt.% ⫽ 1.1113 ⫻ (FeOtotal ⫺ FeOmeasured) wt.%. 3. RESULTS

Six of the nine bulk starting compositions were analyzed by wet chemical methods by ISE Carmichael at U.C. Berkeley. Of the remaining three samples, two (KFS-G and KFS-L) were analyzed with the electron microprobe (details given below). One sample (KFS-D) was too hygroscopic for electron microprobe analysis, and its nominal composition is given instead. The compositions of these nine starting materials (glasses) are reported in Table 1. 2.2. Pt-Loop Equilibration The nine samples were equilibrated at various temperatures in a Deltech furnace with a vertical, alumina muffle tube. A gas-tight, brass mounting cap at the top of the alumina tube held two smaller, vertical ceramic tubes: one held the two wires of the thermocouple, and the other held two thick (1 mm) Pt wire leads. Two 0.5-cm holes on the top and bottom brass caps of the alumina muffle tube provided entry and exit ports for the air with which the samples were equilibrated. Small glass chips of the starting samples were fused onto 0.25-mm Pt wire loops (always in duplicate). The size of the sample beads typically ranged from ⬃2.5 mm to 4 mm in diameter; thus, the surface area of Pt wire in contact with the sample was relatively small. The Pt loops were attached to a Pt cage that was placed at the hot spot of the furnace by suspending it from a thin (0.25 mm) Pt wire, which itself was attached to the two thicker Pt wire leads that fed to the top of the furnace. The samples were equilibrated at temperatures ranging from 1123°C to 1596°C for durations of two to 72.5 h. Temperature was measured with a S-type thermocouple placed adjacent to the samples and corrected by 2°C and 1°C for horizontal and vertical thermal gradients, respectively. The thermocouple was calibrated against the melting temperature of 99.99% gold (1064.43°C; accepted IPTS, 1969) measured three times during the course of the experiments. The melting temperatures of gold were 1065°C at the beginning of the experimental program and 1066°C at the completion of the project. The samples were quenched by running an electrical current through the thick Pt lead wires, causing the thinner wire holding the Pt cage to break. The Pt cage then fell into water at the bottom of the furnace, where a brass cup was attached to the lower end of the alumina muffle tube. Quench rates are estimated to be ⬎100°C/s (Xu and Zhang, 1999).

3.1. Equilibration Times At temperatures near 1600°C, the liquids were equilibrated for no more than 2 h to minimize potassium loss. At such high temperatures, equilibration times of 2 h were found to be sufficient for the attainment of ferric-ferrous equilibrium in four of the nine samples (A, E, G, H). (The criteria for evaluating the achievement of equilibrium are discussed at length below.) For samples B, C, D, F, and L, data at ⬃1600°C could not be obtained because equilibration times longer than 2 h were required, yet led to such severe potassium loss that the liquids entered an immiscibility field and were discarded. At low temperatures near 1100°C to 1200°C, the minimum time required for the attainment of ferric-ferrous equilibrium was 19 to 25 h for the more viscous liquids (samples E, G, and L with ⬎65 wt.% SiO2 and/or ⬍20 wt.% K2O), whereas less than 5 h was required for the less viscous liquids (samples A, B, C, D, F, and H with ⱕ50 wt.% SiO2, and/or ⬎30 wt.% K2O). No data were obtained at ⬃1100°C for sample L, because it is below its liquidus at that temperature. 3.2. Negligible Iron Loss From the microprobe analyses in Table 2, it is clear that negligible iron was lost from the sample melts to the Pt wire during the experiments. This can be attributed to the very low ferrous iron concentrations present in these potassic liquids equilibrated in air. Minimal iron loss to the Pt loop wires was also found during equilibration of Na2O-FeO-Fe2O3-SiO2 and CaO-FeO-Fe2O3-SiO2 liquids in air (Lange and Carmichael, 1989; Kress and Carmichael, 1989, respectively), all of which contained similar or higher FeO concentrations than the liquids in this study.

2.3. Chemical Analyses of Run Products Quenched samples were analyzed for SiO2, K2O, and total iron (as Fe2O3) with a Cameca CAMEBAX electron microprobe. The standards used were a potassium feldspar (GKFS) for K, an orthopyroxene (FESI) for Fe, and diopside (PX69) for Si. An accelerating potential of 30 kV and a sample current of 3 nA were optimal for minimizing alkali loss. Analyses were performed by using point mode and integration times of 30 s for Fe and Si and 10 s for K with 15 to 30 spots per analysis. For the majority of samples, scans of K count rate vs. time indicated that 10-s counting times for K were appropriate for these conditions. The K2O concentrations, for samples in which K-migration occurred during the 10-s interval, were determined by using a time-zero

3.3. Significant Potassium Loss For those samples equilibrated for long durations at high temperatures, loss of potassium was so severe that the liquids moved out of the peralkaline region of the K2O-FeO-Fe2O3SiO2 quaternary system (Fig. 2). A large subset of these hypoalkaline glasses were found to be magnetic, despite the absence of any optically visible quench crystals. Scanning transmission electron spectroscopy images were obtained on these magnetic glasses and revealed the presence of tiny (50 –

Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts

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Fig. 1. Plots of counts per second vs. time for K, Fe, and Si in the standards, GKFS (K-feldspar), FESI (orthopyroxene), and PX69 (diopside), respectively. Counts per second vs. time are also shown for K in KFS-7, KFS-E, and KFE-B (samples in which a time-zero intercept method was used to determine K concentration).

200 nm in diameter), crystallized spherules. The presence of the spherules are interpreted to represent immiscibility between an iron-rich, silica-poor liquid (the spherules) and a silica-rich, iron-poor liquid (the quenched matrix glass). Not all of the magnetic glasses contained spherules; in some, only individual magnetite crystals (25–75 nm) were present. Scanning transmission electron microscopy images were also obtained on four nonmagnetic glasses and were found to be completely homogenous, without spherules or crystals. None of the magnetic glass samples were included in the final data set, and as a precaution, only peralkaline liquids were included. One of the most surprising results is the lack of K2O gradients in the quenched glass beads, despite the loss of substantial

amounts of K2O during the high-temperature equilibrations. For example, sample A lost 28 wt.% K2O (64% relative) during its highest temperature equilibration near 1600°C, and all samples lost significant K2O even at lower equilibration temperatures (Table 2). However, electron microprobe transects across the diameter of the glass beads show chemical homogeneity in SiO2, total iron, and K2O. This rather remarkable feature strongly suggests that the rate of K diffusion within the melt is faster than the rate of K volatilization at the melt/gas interface, thus preventing K2O gradients from developing. The inference that alkali diffusion in the silicate melts is faster than alkali volatilization makes sense when one considers that the diffusion of Na is generally faster than that of K in multi-component

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J. A. Tangeman, R. Lange, and L. Forman Table 2. Experimental conditions and compositions of run products.

Sample

Time (h)

Temp. (K)

SiO2 (wt.%)

SiO2 S. D.

K2O (wt.%)

K2O S. D.

Fe2O3 (wt.%)

Fe2O3 S. D.

FeO (wt.%)

Total

FeO (calc)a

Residual

A-1100 A-1200 A-1200 A-1250 A-1250 A-1300 A-1350 A-1450 A-1450 A-1500 A-1600

8 5 5 5 5 5 5 5 5 5 3

1396 1489 1493 1536 1536 1579 1627 1718 1715 1771 1863

47.44 52.98 52.73 48.30 50.54 60.76 53.74 59.85 61.88 61.43 63.92

0.26 0.46 0.35 0.55 0.35 0.48 0.35 0.22 0.40 0.34 0.38

36.32 29.43 29.28 34.14 32.76 21.20 29.27 21.35 18.23 18.55 15.46

0.29 0.32 0.35 0.27 0.15 0.18 0.30 0.23 0.22 0.21 0.21

15.56 17.57 16.76 16.86 16.46 18.08 16.68 16.98 17.95 17.93 16.70

0.21 0.43 0.32 0.27 0.30 0.21 0.70 0.21 0.18 0.21 0.26

0.27 0.50 0.59 0.52 0.69 0.91 1.09 1.86 1.98 2.35 3.60

99.59 100.48 99.36 99.82 100.45 100.95 100.78 100.04 100.04 100.26 99.68

0.22 0.48 0.47 0.71 0.67 0.83 1.19 1.89 1.90 2.62 3.79

0.05 0.02 0.12 ⫺0.19 0.02 0.08 ⫺0.10 ⫺0.03 0.08 ⫺0.27 ⫺0.19

B-1100 B-1100 B-1100 B-1100 B-1200 B-1200 B-1200 B-1300 B-1400

10 8 5 5 8 5 5 5 5

1401 1396 1407 1400 1495 1491 1493 1581 1674

48.00 41.48 43.52 42.69 48.29 44.76 44.53 46.60 49.14

0.22 0.68 0.49 0.45 0.18 0.36 0.26 0.32 0.61

20.51 30.50 27.92 30.42 20.49 25.81 26.58 23.31 17.90

0.21 0.28 0.29 0.60 0.26 0.18 0.17 0.18 0.43

30.64 27.11 27.34 26.76 29.99 28.14 27.62 28.57 29.27

0.22 0.48 0.26 0.52 0.24 0.33 0.27 0.22 0.77

0.55 0.39 0.55 0.35 0.92 0.68 0.68 1.42 2.93

99.70 99.48 99.33 100.22 99.69 99.39 99.41 99.90 99.24

0.41 0.41 0.43 0.41 0.86 0.84 0.85 1.57 2.74

0.14 ⫺0.02 0.12 ⫺0.06 0.06 ⫺0.16 ⫺0.17 ⫺0.15 0.19

C-1100 C-1150 C-1200 C-1200 C-1200 C-1200 C-1300 C-1300 C-1400 C-1400 C-1400 C-1400 C-1500

5 5 8 5 5 5 5 5 5 5 5 5 5

1400 1447 1495 1484 1492 1497 1584 1588 1678 1674 1674 1674 1764

51.02 54.03 58.06 55.58 56.79 55.12 62.27 57.52 62.87 59.31 61.24 62.32 64.85

0.49 0.50 0.56 0.61 0.43 0.58 0.24 0.51 0.37 0.25 0.39 0.72 0.34

29.65 28.18 22.33 25.30 23.96 25.49 16.33 20.92 16.15 21.59 16.41 15.43 13.03

0.31 0.27 0.24 0.24 0.23 0.77 0.20 0.46 0.19 0.16 0.28 0.61 0.11

18.90 17.82 18.95 18.25 18.50 18.03 20.28 19.59 19.60 17.66 19.66 20.61 18.53

0.45 0.17 0.30 0.38 0.29 0.72 0.16 0.75 0.28 0.20 0.68 0.66 0.18

0.43 0.52 0.65 0.58 0.73 0.82 0.88 0.94 1.52 1.62 1.66 1.47 2.82

100.00 100.55 99.99 99.71 99.98 99.46 99.76 98.97 100.14 100.18 98.97 99.83 99.23

0.26 0.34 0.49 0.45 0.48 0.50 0.92 0.98 1.63 1.53 1.61 1.67 2.45

0.17 0.18 0.16 0.13 0.25 0.32 ⫺0.04 ⫺0.04 ⫺0.11 0.09 0.05 ⫺0.20 0.37

D-1200 D-1300

5 5

1491 1581

34.96 35.42

0.18 0.21

24.67 23.36

0.20 0.22

38.76 38.59

0.25 0.21

0.84 2.15

99.23 99.52

1.31 2.42

⫺0.47 ⫺0.27

E-1100 E-1200 E-1200 E-1300 E-1300 E-1400 E-1400 E-1500 E-1600

30.5 72.5 24 24.5 17 24.5 12.5 4 2

1411 1498 1497 1589 1587 1677 1677 1772 1865

64.49 64.97 64.84 64.27 65.09 66.00 65.75 64.27 65.78

1.04 0.59 1.07 0.58 0.67 0.50 0.45 0.38 0.34

14.36 13.86 13.74 13.48 13.83 12.23 12.61 13.40 12.39

0.41 0.26 0.40 0.27 0.35 0.21 0.15 0.20 0.23

21.01 20.55 20.72 20.72 20.33 19.41 20.14 19.08 17.34

0.69 0.37 0.65 0.50 0.48 0.25 0.31 0.19 0.23

0.06 0.37 0.27 0.88 0.46 1.81 1.33 2.55 3.58

99.92 99.75 99.57 99.35 99.71 99.45 99.83 99.30 99.09

0.25 0.48 0.48 0.94 0.91 1.52 1.59 2.65 3.85

⫺0.19 ⫺0.11 ⫺0.21 ⫺0.06 ⫺0.45 0.29 ⫺0.26 ⫺0.10 ⫺0.27

F-1100 F-1100 F-1200 F-1200 F-1300 F-1300 F-1400 F-1400 F-1500 F-1500

5 5 5 5 5 5 5 5 7 6

1407 1398 1492 1497 1588 1584 1674 1675 1768 1767

51.02 48.56 54.45 53.75 56.25 55.81 60.52 58.88 57.92 59.53

0.45 1.09 0.25 0.39 0.32 0.28 0.43 0.64 0.19 0.40

28.40 28.92 21.70 23.91 19.84 20.81 13.83 17.05 17.61 15.52

0.18 0.70 0.12 0.21 0.15 0.20 0.14 0.27 0.20 0.24

20.85 21.78 22.24 21.71 22.10 22.43 23.15 22.13 21.03 21.29

0.53 1.11 0.14 0.21 0.28 0.20 0.31 0.17 0.08 0.28

0.37 0.42 0.69 0.73 1.15 1.04 2.05 1.93 2.75 3.00

100.64 99.68 99.08 100.10 99.34 100.09 99.55 99.99 99.31 99.34

0.30 0.30 0.58 0.61 1.12 1.12 1.90 1.89 3.09 3.04

0.07 0.12 0.11 0.12 0.03 ⫺0.08 0.15 0.04 ⫺0.34 ⫺0.04

G-1100 G-1300 G-1400 G-1500 G-1500 G-1600 G-1600

24 19 15.5 16.5 4 6 2

1413 1593 1680 1771 1772 1866 1869

67.55 73.32 75.90 85.30 74.91 83.67 77.86

0.54 0.52 0.47 0.43 0.53 0.40 0.47

28.64 23.26 20.35 10.68 20.93 12.31 18.30

0.58 0.34 0.25 0.26 0.38 0.24 0.24

3.43 3.49 3.51 3.78 3.14 3.39 3.06

0.08 0.07 0.05 0.08 0.07 0.09 0.08

0.00 0.12 0.30 0.44 0.49 0.69 0.70

99.62 100.19 100.06 100.20 99.47 100.06 99.92

0.04 0.16 0.27 0.43 0.41 0.66 0.64

⫺0.04 ⫺0.04 0.03 0.01 0.08 0.03 0.06 (Continued)

Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts

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Table 2. Continued Sample

Time (h)

Temp. (K)

SiO2 (wt.%)

SiO2 S. D.

K2O (wt.%)

K2O S. D.

Fe2O3 (wt.%)

Fe2O3 S. D.

FeO (wt.%)

Total

FeO (calc)a

Residual

H-1100 H-1200 H-1250 H-1300 H-1300 H-1350 H-1400 H-1450 H-1500 H-1600

8 5 5 5 5 5 5 5 15 7

1398 1489 1536 1579 1582 1627 1669 1715 1770 1865

56.52 57.72 56.57 63.12 61.67 59.21 71.36 68.70 73.62 75.10

0.56 0.55 0.37 0.38 0.35 0.70 0.37 0.43 0.45 0.43

33.64 31.59 33.31 25.38 29.02 31.13 15.79 20.17 13.38 11.21

0.27 0.52 0.35 0.29 0.26 0.36 0.21 0.14 0.19 0.22

9.36 9.96 9.31 10.47 9.80 9.40 11.69 10.52 11.54 11.06

0.15 0.44 0.19 0.14 0.17 0.27 0.68 0.17 0.11 0.13

0.33 0.44 0.57 0.70 0.70 0.85 0.99 1.38 1.40 2.10

99.85 99.71 99.76 99.67 101.19 100.59 99.83 100.77 99.94 99.47

0.12 0.26 0.35 0.48 0.48 0.64 0.85 1.06 1.47 2.26

0.21 0.18 0.22 0.22 0.22 0.21 0.14 0.32 ⫺0.07 ⫺0.16

L-1200 L-1300 L-1400 L-1500 L-1500

24 19.5 7 6 4

1499 1590 1679 1768 1772

50.41 50.20 50.32 50.52 50.80

0.58 0.48 0.47 0.30 0.23

17.92 16.95 16.80 16.19 15.88

0.29 0.46 0.29 0.34 0.48

30.38 30.27 29.52 27.92 27.29

0.42 0.34 0.52 0.23 0.36

0.73 1.79 2.88 4.77 4.70

99.44 99.21 99.52 99.40 98.67

0.86 1.64 2.80 4.39 4.35

⫺0.13 0.15 0.08 0.38 0.35

a

Calculated using Eqn. 4 and the regression results presented in Table 3.

silicate melts (e.g., Jambon, 1982), and yet less Na is lost by volatilization than K in various high-temperature equilibrations of silicate melts in air (e.g., Kilinc et al., 1983; Lange and Carmichael, 1989). If the rate-limiting step for alkali volatilization is the diffusion of the alkali from the interior of the melt to the gas/melt interface, then more Na would be volatilized than K, which is not observed. The effect of this substantial K volatilization and rapid K diffusion on the iron-redox equilibria is discussed below. 3.4. Evaluation of Ferric-Ferrous Equilibria Although significant potassium volatilization occurred during most experiments, several lines of evidence (Kress and Carmichael, 1988; Lange and Carmichael, 1989) were used to evaluate whether ferric-ferrous equilibrium was achieved (and

maintained) by the time the samples were quenched. First, replicate wet chemical measurements of FeO on splits obtained from the same quenched bead that were within 2␴ error (⫾0.2 wt.%) of each other were taken to indicate homogeneity in each bead with respect to FeO concentration. Second, sectioned glass beads that were homogeneous, on the basis of electron microprobe analyses of SiO2, K2O, and total Fe2O3, were considered to be consistent with the achievement of equilibrium. All those samples that met these two criteria (the 75 liquids in Table 2) were included in our final data set. The achievement of ferric-ferrous equilibrium in our experimental liquids, despite the continual and extensive loss of potassium can be understood from the experiments of Cooper et al. (1996), which were designed to constrain the mechanism and dynamics of oxidation in basaltic melts. It is well known that for melts and glasses of similar polymerization and at similar temperatures, the kinetics of oxygen tracer diffusion are 10 to 1000 times slower than the kinetics of chemical diffusion (e.g., Cook et al., 1990). Cooper et al. (1996) showed, however, that dissipation of a redox gradient does not involve the movement of slower oxygen atoms, but rather is achieved with the more rapid, charge-coupled motion of fast-moving, networkmodifying cations and an even faster counterflux of electron holes. Cooper et al. (1996) concluded that the chemical diffusion response to an oxidation potential is unequivocally dominated by the rapid diffusion of network-modifying cations, and not by the slower diffusion of oxygen species. Therefore, maintenance of ferric-ferrous equilibrium in our experimental charges is expected despite the severe loss of potassium, because of the following differences in rates: electron hole migration in the melt ⬎ K diffusion in the melt ⬎ K volatilization at the melt/gas interface. 3.5. Variation in the Redox Ratio with Temperature and Composition

Fig. 2. A plot of run products in the K2O-SiO2-Fe2O3 ternary system (mol.%; projected from the FeO component). The peralkaline-hypoalkaline boundary is marked by the solid line.

The data presented in Table 2 demonstrate an increase in the ferrous iron concentration in all melts as a function of temperature. This trend is consistent with the following reaction:

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J. A. Tangeman, R. Lange, and L. Forman

Fig. 3. Plots of Fe3⫹/FeT vs. K2O and vs. total iron for the K2O-FeO-Fe2O3-SiO2 (KFS) liquids from this study. Plots of Fe3⫹/FeT vs. Na2O and vs. total iron from the data of Lange and Carmichael (1989) on Na2O-FeO-Fe2O3-SiO2 (NFS) liquids. Plots of Fe3⫹/FeT vs. CaO and vs. total iron from the data of Kress and Carmichael (1989) on CaO-FeO-Fe2O3-SiO2 (CFS) liquids.

FeO1.5(liq) ⫽ FeO(liq) ⫹ 0.25 O2(gas)

(1)

in which temperature favors the higher entropy side of the reaction. The effect of composition on the direction of this reaction is explored by plotting Fe3⫹/FeT vs. K2O and FeOT (total iron as FeO) concentration at 1100, 1200, 1300, 1400, and 1500°C in Figure 3. The data indicate that the proportion of ferric iron systematically decreases with increasing K2O at constant temperature. The Fe3⫹/FeT trend is less clear, with more scatter, as a function of total iron content. Similar plots are presented in Figure 3 for the iron-redox measurements on Na2O-FeO-Fe2O3-SiO2 liquids (Lange and Carmichael, 1989) and on CaO-FeO-Fe2O3-SiO2 liquids (Kress and Carmichael, 1989). The combined results indicate that liquids of similar

bulk composition, but with K2O substituted for Na2O (or Na2O substituted for CaO), have higher Fe3⫹/FeT ratios, similar to the trend found for magmatic liquids. However, within each quaternary system, the variations in Fe3⫹/FeT with composition are completely different. In calcic liquids, the iron-redox ratio is nearly independent of the total iron content and is primarily controlled by the CaO concentration, with the ferric-ferrous ratio increasing with CaO (Kress and Carmichael, 1989). In sodic liquids, the iron-redox ratio is strongly correlated to, and increases with, the total iron content (Lange and Carmichael, 1989). The potassic system is similar to the calcic system in that the iron-redox ratio is less dependent on the total iron content and more strongly con-

Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts

1815

trolled by the K2O concentration. However, in direct contrast to the calcic system, the ferric iron concentration decreases with increasing K2O concentration. These differences in the ironredox trends between the potassic, sodic, and calcic liquids are explored more fully in Section 4. 3.6. Development of a Model Equation The ferric-ferrous equilibrium reaction in Eqn. 1 can be modeled as a function of both temperature and composition with the following equation: 0 ⫽ ⌬H° ⫺ T ⌬S° ⫹ RT ln共a FeO/a FeO1.5) ⫹ 0.25 RT ln f O2

(2)

where ⌬H° and ⌬S° refer to the standard-state enthalpy and entropy of Eqn. 1, a FeO and a FeO1.5 are the activities of the FeO and FeO1.5 liquid components, and f O2 represents the fugacity of oxygen in the system. The temperature dependence to ⌬H° and ⌬S° can be neglected over the temperature interval of these experiments (1115–1587°C), as was found by Lange and Carmichael (1989) and Kress and Carmichael (1989) in their study of ferric-ferrous equilibria in sodic and calcic liquids, respectively. By rearranging Eqn. 2 and applying regular solution theory, the following relationship is attained: ln

X FeO1.5 ⌬H° ⌬S° ⫽ ⫺ ⫹ 0.25 ln f O2 X FeO RT R ⫹

冘

X i共W i,FeO ⫺ W i,FeO1.5兲 RT

y⫽ (3)

where W i, j are Margules binary interaction parameters, which account for the effect of bulk composition on ferric-ferrous equilibria. The parameters in the last term modifying the mole fraction of each oxide component (X i ) do not define the magnitude and sign of individual binary interaction parameters, but rather the relative differences between the two. The form of Eqn. 3 produces a bias if it is used in a linear least-squares regression because of the way the experimental errors propagate during construction of the dependent variable, ln(X FeO1.5/X FeO). This point is illustrated in Figure 4, where calculated values of ln(X FeO1.5/X FeO) with two sigma error bars are plotted vs. inverse temperature (1000/RT) for sample KFS-E, which has an initial total iron concentration of 20.8 wt.% Fe2O3. Because errors in ln(X FeO1.5/X FeO) are significantly larger when FeO values are small (at low temperature), a regression with ln(X FeO1.5/X FeO) as the dependent variable effectively weights the low-temperature experiments far more heavily than those at higher temperatures. However, the analytical error in wt.% FeO is equal at all temperatures and for all values of X FeO. An effective alternative is to simply use the measured parameter, wt.% FeO, as the dependent variable in the regression because the errors in wt.% FeO are uniform. Eqn. 3 can thus be rearranged: wt.% FeO ⫽ MW FeOn FeO1.5 exp(⫺y)

Fig. 4. A plot of ln(X FeO1.5/X FeO) with two sigma error bars vs. 1000/RT for sample KFS-E (initial total iron ⫽ 20.8 wt.%) at several temperatures between 1225°C and 1592°C. The errors increase strongly with decreasing FeO content and therefore with decreasing temperature. The sample has so little FeO (0.06 wt.%) at the lowest temperature of equilibration (1135°C), leading to such a large error on ln(X FeO1.5/X FeO), that it is not plotted. This sample lost relatively little K2O during the high-temperature equilibrations and therefore is a good test of the applicability of the enthalpy obtained from the solid standard-state reaction (⌬H° ⫽ ⬃134 kJ/mol). The straight line reflects a slope of this magnitude and fits the liquid data very well, especially at high temperatures where the values of ln(X FeO1.5/X FeO) have the lowest errors.

(4)

where MW FeO is the molecular weight of FeO, n FeO1.5 is the number of moles of FeO1.5, and

⌬H° ⌬S° ⫺ ⫹ 0.25 ln f O2 ⫹ RT R

冘

X i共W i,FeO ⫺ W i,FeO1.5兲 . RT (5)

where

冘

X i共W i,FeO ⫺ W i,FeO1.5兲 1 ⫺ W SiO2,FeO1.5兲 ⫽ 共X 共W RT RT SiO2 SiO2,FeO

⫹ X K2O共W K2O,FeO ⫺ W K2O,FeO1.5兲 ⫹ W FeO,FeO1.5共X FeO1.5 ⫺ X FeO兲兲. (6) This nonlinear regression equation (Eqns. 4 and 5) was minimized by using a modified Marquardt procedure (Nash, 1979; Algorithm 23) to extract the following unknowns from the data set: ⌬S°, W FeO,FeO1.5 and W i,FeO ⫺ W i,FeO1.5 for i ⫽ K2O and SiO2. Calorimetric data presented in Robie and Hemingway (1995) for stoichiometric crystalline FeO, Fe2O3, and O2 gas were used to estimate a value of 133.9 kJ/mol for ⌬H° at a reference temperature of 1300°C. Because this value for ⌬H° is based on the solid standard-state reaction (because pure FeO and pure FeO1.5 liquid are not stable), the fitted binary interaction parameters include the difference between the molar heats of fusion of the FeO and FeO1.5 components (which is likely a small number) as well as the non-zero heats of mixing. The validity of this value for ⌬H° is demonstrated in Figure 4 in the plot of ln(X FeO1.5/X FeO) vs. 1000/RT for KFS-E. For samples of constant bulk composition (little loss of K2O), the slope of ln(X FeO1.5/X FeO) vs. 1000/RT is a measure of the enthalpy of the liquid iron-redox reaction. KFS-E is the only sample that did not lose substantial amounts of K2O, and therefore is a good test of the imposed enthalpy of ⬃134 kJ/mol. From Figure 4 it can be seen that the data are fit very well, especially at the high temperatures (high FeO contents) where the parameter ln(X FeO1.5/X FeO) is best constrained. The

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J. A. Tangeman, R. Lange, and L. Forman Table 3. Results of the nonlinear regression.a

wt.% FeO ⫽ MW(FeO) ⴱ nFeO1.5 ⌬H° (1300°C) ⌬S° (1300°C) WFeO,SiO2 ⫺ WFeO1.5,SiO2 WFeO,K2O ⫺ WFeO1.5,K2O a b

ⴱ exp(⫺y) 133.9 kJ/molb 62.0 ⫾ 1.3 kJ/mol 䡠 K 14.1 ⫾ 5.4 kJ/mol ⫺10.8 ⫾ 0.4 kJ/mol

Standard error of the fit is 0.18 (2␴). Estimated from calorimetric data (see text).

error bars on this term become excessively large at low temperatures, where there is so little FeO in the liquid that the concentration approaches the error of the measurement. The results of the regression are given in Tables 2 and 3, and demonstrate the excellent quality of the fit. A key feature of the regression is the well-resolved, negative value (⫺10.8 ⫾ 0.4 kJ/mol) for the fitted parameter, W K2O,FeO ⫺ W K2O,FeO1.5, which modifies the mole fraction of K2O in Eqn. 5. This result supports the trend observed for the experimental data at all temperatures in Figure 3, namely that the effect of increasing K2O concentration is to decrease the ferric-ferrous ratio in K2O-FeO-Fe2O3-SiO2 liquids. A second important feature of the regression is that the value for W FeO,FeO1.5 is statistically indistinguishable from zero, and was therefore dropped from the regression. Because this parameter modifies the term, X FeO1.5 ⫺ X FeO, in Eqn. 5, and because X FeO1.5 greatly exceeds X FeO in all experimental liquids, the zero value for W FeO,FeO1.5 suggests that the ferric-ferrous ratio in peralkaline K2O-FeOFe2O3-SiO2 liquids is independent of the total iron concentration. This is consistent with the weaker correlation between Fe3⫹/FeT and total iron (vs. K2O) seen in Figure 3. The resultant model equation given in Table 3 leads to a standard error of ⫾0.18 wt.% in the prediction of FeO. The uncertainty in the model is well within the combined wet chemical and electron microprobe analytical errors. 4. DISCUSSION

4.1. Compositional Controls of Ferric-Ferrous Equilibria One of the most surprising results of this study is the observation that an increase in K2O concentration causes a decrease in the ferric-ferrous ratio in peralkaline K2O-FeO-Fe2O3-SiO2 liquids. Although this trend is similar to that found in ironbearing, peralkaline K2O-Al2O3-SiO2 liquids by Dickenson and Hess (1981), it is opposite to what they observed in their peraluminous liquids. The trend of decreasing ferric iron with increasing K2O is also exactly opposite that observed in magmatic liquids (Fudali, 1965; Sack et al., 1980; Thornber et al., 1980; Kilinc et al., 1983; Kress and Carmichael, 1988) and that observed in K2O-SiO2 liquids with dilute concentrations (0.4 wt.%) of total iron (Paul and Douglas, 1965). However, the results of our study, as well as those of the studies cited above, can be understood in the context of three, different (though related) compositional effects on ferric-ferrous equilibrium in multicomponent silicate melts. The first effect is the charge-balancing role of alkali and alkaline-earth cations with respect to ferric iron in tetrahedral coordination. Charge-balancing increases the stability of tetrahedrally coordinated ferric iron ([4]Fe3⫹) in magmatic liquids,

in the order: K⫹ ⬎ Na⫹ ⬎ Ca2⫹ (Fudali, 1965; Sack et al., 1980; Thornber et al., 1980; Kilinc et al., 1983; Kress and Carmichael, 1988). Fe2⫹ may also play a charge-balancing role (Dickenson and Hess, 1981; Mysen et al., 1984). It is this increased stability of Fe3⫹ by charge-balancing that leads to an increase in the ferric-ferrous ratio with increasing alkali concentration. The second control involves composition-induced change in the average coordination of ferric iron (known to occur as [4] Fe3⫹, [5]Fe3⫹, and [6]Fe3⫹ in various silicate minerals and glasses). Because the ferric-ferrous ratio increases (at constant total iron content) when Fe3⫹ is in tetrahedral vs. octahedral coordination (Baiocchi et al., 1982; Mysen, 1987; Hannoyer et al., 1992), a composition-induced change in the average coordination number of Fe3⫹ is anticipated to affect the ferricferrous ratio at constant temperature, total iron, and f O2. It is this second control that may explain the difference between the results of this study on iron-rich K2O-SiO2 liquids and those of Paul and Douglas (1965) on iron-poor K2O-SiO2 liquids (described in detail in a later section). The third compositional control on the iron-redox equilibrium is related to the basicity of melts, namely the activity of the O2⫺ component in silicate melts, as seen in the following homogeneous liquid reaction: Fe2⫹ ⫹ 1/4 O2 ⫽ Fe3⫹ ⫹ 1/2 O2⫺.

(7)

The basicity of a melt is a measure of the degree of negative charge on the oxygen atoms and, hence, their electron donor power (Duffy, 1996). It is possible to quantify this parameter in a glass by measuring its optical basicity (⌳), using orbital expansion spectroscopy in the ultraviolet region. As explained by Duffy (1989), not only is optical basicity a measurable quantity, but it can also be modeled as a linear function of the individual oxide components. From this work, it has been established that at equivalent bulk compositions, the basicity of a multi-component silicate glass is increased in the order K ⬎ Na ⬎ Ca. Moreover, along the R-SiO2 join (R ⫽ Na2O, K2O, CaO), increasing R increases the basicity of the glass. Therefore, the reaction in Eqn. 7 predicts that at constant f O2 (in air), an increase in the activity of O2⫺ (achieved by increasing K2O in the K2O-FeO-Fe2O3-SiO2 system) will lead to a decrease in the ferric-ferrous ratio. This is exactly the trend observed for the liquids in this study (Fig. 3). 4.2. Magmatic vs. K2O-FeO-Fe2O3-SiO2 Liquids The discrepancy between the effect of K2O on the ferricferrous ratio in magmatic liquids vs. those of this study may be related to the nonperalkaline nature of the experimental magmatic liquids (Fudali, 1965; Sack et al., 1980; Thornber et al., 1980; Kilinc et al., 1983; Kress and Carmichael, 1988), enhanced by rapid alkali loss during the high-temperature equilibrations. In those liquids, the concentration of the alkalis was commonly much less than that of alumina and ferric iron combined. Therefore, for any increase in K, there was an increase in the concentration of Fe3⫹ being charge-balanced and thus a concomitant increase in the ferric-ferrous ratio. In contrast, the liquids examined in this study are predominantly peralkaline. Because K2O ⱖ Fe2O3, an increase in the K2O

Ferric-ferrous equilibria in K2O-FeO-Fe2O3-SiO2 melts

concentration in the melt cannot increase the abundance of Fe3⫹ being charge-balanced by K⫹. Therefore, the chargebalancing effect of K⫹ on Fe3⫹ is not expected to be a primary control on the iron-redox ratio for the liquids in this study. This explanation is supported by the results of Dickenson and Hess (1981) on iron-bearing K2O-Al2O3-SiO2 liquids. They found that the ferric-ferrous ratio increases dramatically with K2O in peraluminous liquids, but decreases slightly with increasing K2O in peralkaline liquids. When CaO is added to the system, the results are broadly similar (Dickenson and Hess, 1986). Dickenson and Hess (1981, 1986) argue that increases in the average coordination of both Al3⫹ and Fe3⫹ in the peraluminous region is the principal cause of the dramatic effect of increasing K2O on increasing ferric-ferrous ratios, not seen in the peralkaline region. 4.3. The Coordination of Fe3ⴙ in Silicate Melts and Glasses Discussion of the possible role of composition-induced coordination change on ferric-ferrous equilibria requires a brief review of what is known about the coordination of Fe3⫹ in silicates, a topic of considerable uncertainty. Hannoyer et al. (1992) provide an extensive list of published, estimated coordination numbers for Fe3⫹ in alkali silicate and soda lime silicate glasses, determined by using a variety of techniques including optical, electron paramagnetic resonance, and Mo¨ssbauer spectroscopies, XANES spectra and extended XAFS spectroscopy. Most of these studies conclude that ferric iron is predominantly in tetrahedral coordination in alkali and soda lime silicate glasses (Steele and Douglas, 1965; Kurkjian and Sigety, 1968; Hirao et al., 1979; Fox et al., 1982; Greaves et al., 1984; Brown et al., 1986; Wang et al., 1993, 1995). However, some studies conclude that both tetrahedral and octahedral Fe3⫹ are present in alkali silicate and soda lime silicate glasses (Levy et al., 1976; de Grave, 1980; Fenstermacher, 1980; Calas and Petiau, 1983; Wang and Chen, 1987; Hannoyer et al., 1992). There is, in addition, Mo¨ssbauer evidence that Fe3⫹ may be present in pentacoordinated sites in some silicate glasses and melts, including those of acmite composition (Bychkov et al., 1993). Although it is difficult to distinguish between a single population of [5]Fe3⫹ vs. a combined population of [4]Fe3⫹ and [6] Fe3⫹ from Mo¨ssbauer spectra, [5]Fe3⫹ has been found in a few silicate minerals, including vesuvianite (Olesch, 1979), andalusite (Abs–Wurmbach et al. 1981), and orthoericssonite (Halenius, 1995). It is therefore quite likely that ferric iron exists in all three coordination environments in silicate melts and glasses, and that the average coordination number (the relative abundance of [4]Fe3⫹, [5]Fe3⫹, and [6]Fe3⫹) varies with melt composition. In a Mo¨ssbauer study of iron-bearing calcic silicate glasses, Kurkjian and Sigety (1968) showed that the concentration of [4] Fe3⫹ increases relative to [6]Fe3⫹ with increasing CaO/SiO2 ratio. Because the ferric-ferrous ratio tends to increase (at constant FeOtotal) when Fe3⫹ is in tetrahedral vs. octahedral coordination (Baiocchi et al., 1982; Mysen, 1987; Hannoyer et al., 1992), this shift toward Fe3⫹ lower coordination may explain the increase in ferric iron with increasing CaO content as observed by Kress and Carmichael (1989) and seen in Figure 3.

1817

This dependence of the ferric-ferrous ratio on the coordination of Fe3⫹ is similar to that documented in K2O-FeO-Fe2O3P2O5 liquids of composition: xK2O-(100 ⫺ x)P2O5-7Fe2O3, where x ⬎ 30 (Nishida et al., 1981). In that combined Mo¨ssbauer and ESR (electron spin resonance) spectroscopic study, it was found that with increasing K2O/P2O5 (when x ⫽ 30–70), the average coordination of Fe3⫹ systematically decreases ([6]Fe3⫹ shifts toward [4]Fe3⫹), whereas the ferric-ferrous ratio increases (exactly as observed for the calcic silicate liquids). In contrast, when x (K2O concentration) ranges between 0 and 30, the coordination of Fe3⫹ is sixfold and unchanging. For these liquids, the ferric-ferrous ratio decreases with increasing K2O/ P2O5, similar to that observed in this study for the potassic silicate liquids. In other words, when there is no change in the coordination of Fe3⫹ with increasing K2O, the ferric-ferrous ratio appears to be controlled by the reaction in Eqn. 7, namely it increases with the activity of O2⫺ (Nishida et al., 1981). Thus, the iron-redox pattern observed in this study may be explained if the coordination of Fe3⫹ is a constant throughout the compositional range of the K2O-FeO-Fe2O3-SiO2 liquids. This is supported by liquid density measurements of Lange et al. (1998) on samples KFS-A, -B, -C, -E, and -F, which show that the partial molar volume of the Fe2O3 component, V៮ Fe2O3, (highly sensitive to variations in Fe3⫹ coordination; Dingwell and Brearley, 1988) is a constant in these six liquids and thus independent of composition. In contrast, Lange et al. (1998) found that values of V៮ Fe2O3 increase with increasing total iron content in Na2O-FeO-Fe2O3-SiO2 liquids, suggesting that ferric-ferrous ratios in this system vary with changes in the proportion of tetrahedral Fe3⫹ in the melt. 4.4. Dilute vs. Iron-Rich K2O-SiO2 Liquids Another series of iron-redox experiments to consider are those of Paul and Douglas (1965) on extremely dilute (⬍0.5 wt.% Fe) Li-, Na-, and K-silicate melts. They showed that the effect of increasing alkali concentration (e.g., 10 –30 mol.% K2O) was to increase the ferric-ferrous ratio, exactly opposite the trend observed in this study. The explanation for this discrepancy may reside in the difference in total iron concentration between the dilute samples of Paul and Douglas (1965) and the iron-rich samples of this study (Table 2). Paul and Douglas (1965) showed that at dilute iron concentrations (0.04 – 0.50 wt.% Fe), the ferric-ferrous ratio in alkali silicate liquids (e.g., 30 wt.% K2O-70 wt.% SiO2) sharply increases with increasing iron content, then gradually begins to flatten and become constant at ⬎1 wt.% Fe. Although no mechanism was proposed to explain this behavior, one possibility is that at very low iron concentrations (e.g., 0.04 wt.% Fe), Fe3⫹ may reside predominantly in octahedral coordination in alkali silicate melts, and with increasing iron content (0.04 – 1.0 wt.% Fe) may gradually shift to lower coordination. This is supported by the EXAFS measurements of Park and Chen (1982), which show that Fe3⫹ is predominantly in octahedral coordination when present in dilute concentrations in sodium disilicate glasses. It thus appears that the iron-redox results of Paul and Douglas (1965) on a variety of K2O-SiO2 liquids with only 0.4 wt.% total Fe were performed on liquids in which some of the ferric iron was in octahedral coordination. If so, then increasing the K2O/SiO2 ratio is expected, by analogy with

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J. A. Tangeman, R. Lange, and L. Forman

the calcic silicate liquids and potassic phosphate liquids, to promote a shift from [6]Fe3⫹ to [4]Fe3⫹, and thereby increase the ferric-ferrous ratio. This explanation is untested, but could be resolved by a series of Mo¨ssbauer and EXAFS spectroscopic studies on iron-bearing potassic silicate glasses. 5. SUMMARY

A variety of compositional features control the iron redox equilibria in multi-component silicate liquids, including charge-balancing effects on Fe3⫹, changes in the coordination of Fe3⫹ (and possibly Fe2⫹), and the activity of O2⫺ in the melts. It is likely that all three of these interrelated effects play a role in determining the compositional dependence of the ferric-ferrous ratio in magmatic liquids. In contrast, for the sodic and calcic, four-component liquids, we suggest that Fe3⫹ coordination change plays the single most important role in controlling their redox compositional trends. However, for the series of peralkaline K2O-FeO-Fe2O3-SiO2 liquids examined in this study, where neither charge-balancing effects nor variations in Fe3⫹ coordination appear to be major factors, melt basicity is the dominant control on ferric-ferrous equilibria, as expressed in Eqn. 7. The effect of increasing K2O concentration is to increase the activity of O2⫺ and drive the reaction to the right, reducing the ferric-ferrous ratio. This explains the rather surprising result to petrologists that increasing K can reduce the ferric iron content in these simple liquids. Acknowledgments—This study was supported by National Science Foundation Grants EAR-95-08133 and EAR-9706075 to R. Lange. We are especially grateful to Ian Carmichael for performing wet chemical analyses on six of the nine KFS glasses. Carl Henderson gave invaluable assistance with electron microprobe (funded by EAR 82-12764). Two anonymous reviewers and the editor, M. Ghiorso, contributed greatly to the improvement of this manuscript. We also thank B. Mysen, V. Kress, and E. Essene for constructive and helpful comments on an earlier draft. Associate editor: M. S. Ghiorso REFERENCES Abs–Wrumbach I., Langer K., Seifert F., and Tillmanns E. (1981) The crystal chemistry of (Mn3⫹, Fe3⫹)-substituted andalusites (viridines 3⫹ and kanonaite), (Al1⫺x⫺yMn3⫹ x Fey )3-(O兩SiO4): Crystal structure refinements, Mo¨ssbauer and polarized optical absorption spectra. Z. Kristallogr. 155, 81–113. Baiocchi E., Bettinelli M., Montenero A., and di Sipio L. (1982) Spectroscopic behavior of iron(III) in silicate glass. J. Am. Ceram. Soc. 65, 39 – 40. Brown G. E. Jr., Waychunas G. A., Ponader C. W., Jackson W. E., and McKeown D. A. (1986) EXAFS and NEXAFS studies of cation environments in oxide glasses. J. Phys. Colloq. (Paris) 47, 661– 668. Bychkov A. M., Borisov A. A., Khramov D. A., and Urusov V. S. (1993) Change in the immediate environment of Fe atoms during the melting of minerals (Review). Geochem. Int. 30, 1–25. Calas G. and Petiau J. (1983) Coordination of iron in oxide glasses through high-resolution K-edge spectra: Information from the preedge. Solid State Commun. 48, 625– 629. Carmichael I. S. E. (1991) The redox states of basic and silicic magmas: A reflection of their source regions? Contrib. Mineral. Petrol. 15, 24 – 66. Cook G. B., Cooper R. F., and Wu T. (1990) Chemical diffusion and crystalline nucleation during oxidation of ferrous iron-bearing magnesium aluminosilicate glass. J. Non-Cryst. Solids. 120, 207–222. Cooper R. F., Fanselow J. B., Weber J. K. R., Merkley D. R., and Poker D. B. (1996) Dynamics of oxidation of Fe2⫹-bearing aluminosilicate (basaltic) melt. Science 274, 1173–1176.

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