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A mass spectrometric and vibrational spectroscopic investigationof the speciation of water in volcanic glasses and some hydrousminerals

Pandya, Naresh, Ph.D.

University of Hawaii, 1992

V·M·I300 N. Zeeb Rd.Ann Arbor.MI 48106

A MASS SPECTROMETRIC AND VIBRATIONAL SPECTROSCOPIC

INVESTIGATION OF THE SPECIATION OF WATER IN VOLCANIC

GLASSES AND SO:rv1E HYDROUS MINERALS

A DISSERTATION SUBMITTED TO THE GRADUATE DIVISION OF THEUNIVERSITY OF HAWAII IN PARTIAL FULFILLMENT

OF THE REQUIRE:rv1ENTS FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

IN CHEMISTRY

AUGUST 1992

By

Naresh Pandya

DISSERTATION COMMITTEE:

David W. Muenow, ChairmanShiv K. Sharma

Karl SeffRichard G. InskeepMichael O. Garcia

Dedicated

To my motherand late father

iii

iv

ACKNOWLEDGEMENTS

I would like to express my sincere thanks to my research advisor Dr. David

Muenow who has always been a constant source of guidance and encouragement. The

spectroscopic portion of my research has been under the guidance of Dr. Shiv Sharma,

who was never at a loss for ideas. Special thanks go to Jo Ann Sinton for her help in the

preparation of thin-sections. A special mahalo to Drs. Tom Cooney and S. Wang who were

a constant source of help. The time and effort put in by my other committee members: Drs.

Karl Seff, Michael Garcia and Richard Inskeep are greatly appreciated. I would also like to

thank all the support staff in the Chemisty department and Hawaii Institute of Geophysics

for their help. The work carried out in the Raman and infrared spectroscopy laboratory was

supported in part by a research grant from the National Science Foundation to Dr. Shiv

Sharma (EAR-8915830). Finally, financial support provided by the Department of

Chemistry in the form of a teaching assistantship is gratefully acknowledged.

v

ABSTRACT

The first portion of this work employed a combination of high-temperature mass

spectrometry (HTMS) and Fourier-transform infrared (FTIR) spectroscopy to determine

the speciation of water within submarine volcanic glasses. Glasses ranged in composition

from basalts to dacites and are from a variety of tectonic settings. HTMS provided total

water contents and FTIR the molecular water-to-total water ratio. A molar absorptivity of

61 ± 1 L/mol-cm was determined for the fundamental OH-stretching band at 3550 cm- 1 in

these glasses. It was found that the relative abundances of molecular water and hydroxyl

groups depend not only upon total water but also bulk silicate chemistry. For glasses with

total H20 contents between 0.5 and 1.0 wt. %, the ratio of molecular to total water varies

approximately lO-fold from 0.03 to 0.30. Increasing the silica content and/or decreasing

the concentration of non-tetrahedral cations enhances the relative abundances of molecular

water. It was also found that the more slowly cooled glasses have a greater proportion of

molecular water. NMR, IR and Raman results for synthetic alkali silicate glasses hydrated

at ambient conditions show no evidence for hydroxyl groups.

In the second part of this study, phlogopite, biotite and muscovite micas were

studied using FTIR and Raman spectroscopy. In the low frequency region (130-800 cm-1)

at ambient conditions the Raman spectra are very similar. Muscovite has a unique band at

-266 cm-1 which is attributed to the formation of a O-H-O isosceles triangle. The position

of the hydroxyl stretching band changes significantly between phlogopites and muscovites.

At high pressure (- 200 kbar), the hydroxyl bands for muscovite and phlogopite were

shifted in opposite directions with increasing pressure, and is attributed to differences in

hydroxyl orientation. From high (-675K) and low (-33K) temperature studies it was found

that the band at -1635 cm-1 is attributed to combination bands of the aluminosilicate

network and not to adsorbed water, whereas the broad hydroxyl band (-3625 em -1) for

muscovite is due to the presence of static disorder.

viTABLE OF CONTENTS

ACKNOWLEDGEl\1ENTS.................................................................. ivABSTRACT v

LIST OF TABLES ,. viii

LIST OF FIGURES , , . . . .. . ixPART 1CHAPTER I. WATER IN SILICATE MELTS.................................. 1

A. Introduction 1B. Mechanisms for H20 solubility................................................. 3C. Mechanisms for C02 solubility 14D. The mass spectrometriclFTIR study........................................... 15

CHAPTER II. EXPERIl\1ENTAL l\1ElHODS......... 17A. High temperature mass spectrometry '" . .. 17B. Infrared spectroscopy , . . . .. . .. . .. 18

1. Molecular explanation 202. Vibrationalmodes , 213. Instrumentation... 22

C. Sample preparation , , 251. Mass spectrometry. . . .. . . .. . . . . . .. . . .. . . .. .. .. . . . . . . .. . . .. . . .. . . . . . . .. . .. 252. Infrared....... 28

D. Instrumental methods , 28E. Density measurement. , " . . . . . . . .. . .. 29F. Magic angle spinning nuclear magnetic resonance. .. .. 30

CHAPTER III. RESULTS AND DISCUSSION................................. 33A. Results........... .. . .. . . . . . .. 33B. Discussion.......... 49

1. Effect of bulk composition on the speciation of water. . . . . . . . . . . . . .. 492. Effect of quenching rate and annealing " 533. Comparison with previous studies............. 56

C. Future work '" , . '" . . . 58

APPENDIX A: Obsidian (An interlaboratory comparison) , . 59

APPENDIX B: The speciation of water within hydrous alkali silicate glasses 64

APPENDIX C: The effect of quenching rate and temperature on themolecular-to-total water ratio................................................. 78

1. Depth profile study of the glassy rind from Loihi seamount........ 782. The effect of annealing upon the ratio of molecular-to-total

water within a volcanic glass.......................................... 86

APPENDIX D: Water poor and water rich glasses........................................ 911. Water poor glasses 912. Water rich glasses. . . . . . . . .. . . . . . . . . . . . . .. .. .. . .. . . . .. . . . . . . .. . . . . . . . . . .. 95

REFERENCES FOR PART 1 100

viiTABLE OF CONTENTS (continued)

PART~.

CHAPTER IV. SPECTROSCOPIC STUDY OF MICAS ATAMBIENT CONDITIONS...................................... 108

A. Structure and composition....................................................... 108B. Study of micas.................................................................... 111C. Infrared spectroscopy............................................................ 113D. Raman spectroscopy 119

CHAPTER V. EXPERIMENTALMElliODS 121A. Infrared spectroscopy.... 121B. Raman spectroscopy............................................................. 121

1. Raman scattering......................................................... 1232. Stokes and anti-Stokes lines............................................ 1233. Classical mechanics of Raman scattering.............................. 1284. Instrumentation........................................................... 1305. Sources.................................................................... 1306. Monochromators . .. . . . . . . .. . ... . . .. . . . . . . . . . . . . ... . . .. . . .. .. . .. .... . .. ... 1317. Multichannel Raman spectroscopy: A triplemate spectrograph.. . .. 131

C. Instrumental methods.... 132

CHAPTER VI. RESULTS AND DISCUSSION......................................... 136A. Samples............................................................................ 136B. Raman low wavenumbers region (130-800 em-I) 139C. Infrared region (1600-2000 em-I)............................................. 154D. Raman and infrared hydroxyl stretchingregion

(3000-4000 em -1) 155E. Near infrared region (4000-4500 em-I)........................................ 156

APPENDIX E: The effect of high pressure on the infrared spectra of micas.......... 1571. Experimental... 1582. Results 1633. Discussion... 168

APPENDIX F: The effect of high temperature on the IR spectra of micas............. 1751. Experimental........................................................ 1752. Results and discussion 176

APPENDIX G: The effect of low temperature on the Raman spectra of muscovitein the hydroxyl stretchingregion................ 182

1. Experimental.............................. 1822. Results and discussion 183

APPENDIXH:1. High temperature mass spectrometry (HTMS) 187

(i). Experimental................................................... 189(ii). Results and discussion.. 189

2. Electron microprobe analyses........................................... 192

REFERENCES FOR PART 2............................................................... 193

viii

LIST OF TABLES

TABLE

1 Major elements and H20 abundances forvolcanic glasses (wt.%) 40

2 Data for fundamental OH stretch peak (3550 cm- l)..................... 42

3 Data for fundamental H20 bending peak (1635 cm- l ) 43

4 Data for low water, high silica glasses............... . 44

5 Near-infrared data for glasses from the eruption (ca. 1340 AD)of the Mono Craters chain in central California................. ..... .... 61

6A Loihi depth profile: volatile abundances (wt.%)......................... 79

6B Microprobe analysis of Loihi 1-4 glass(wt.%) 79

7 The ratio I(l636)/I(3550) for the depth profile study.... 84

8 The re~ati0!1shipbetween the I(l636)/I(3550) ratio withannealing orne. . . .. . . . . . . . . . . . . . .. . .. . . .. . . .. . . .. . . . . . .. . . .. . . .. . .. . ... . . .... . 90

9 Major elemental and infrared data for low water content glasses...... 92

10 Major elements, volatiles and infrared data for waterrich glasses................................................................... 97

11 Comparison of extinction coefficient values obtained in this studywith those of other workers. . . . . . .. . . .. . 99

12 Chemical analyses ofa typical mica (wt.%).............................. 137

13 Band positions and assignments for the Raman andinfrared spectra of micas. . . . . . . . . . . .. .. .. . . .. .. . . . . . . . . .. . . .. . . .. . . .. . . . .. .. 140

14 The effect of pressure on the position and FWHM of thehydroxyl band for muscovite 164

15 The effect of pressure on the position and FWHM of thehydroxyl band for phlogopite (K8) 166

16 Effect of low temperature on the OH band position ofmuscovite........ 184

17 Elemental analyses for a muscovite and biotite sample (wt.%)........ 191

FIGURE

1

2

3

4

5

6

7

8

9

10

11

12

13

14A

LIST OF FIGURES

The system albite-anorthite-water at high and low pressures .

The solubility of water in silicate melts , .

Summary of published OHJH20 distribution in silicate glassesas a function of total water concentration .

Schematic of the high-temperature quadrupole mass spectrometerfacility .

Modes of vibration of the H20 molecule .

Weight percent H20 as a function of ocean depth .

Mass pyrograms showing the release profile of waterfor (a) altered glass (b) unaltered glass .

MAS NMR Qn notation for silicate anions .

Mass pyrogram for a Mahukona (MA-21) submarine basaltic glassshowing the degassing behavior of H20 (rn/e = 18) with respect

to temperature .

Mass pyrograms for a Juan de Fuca (JDF-21) basaltic glassshowing the degassing behavior of H20, C02, and S02with temperature .

Infrared absorption spectrum of a typical submarine volcanicglass between 1400-4000 em-I .

Infrared spectra for Mariana arc volcanic glasses in the

4000-1200 cm- 1 region .

Relationship between the band intensity for the fundamentalOH-stretch at -3550 cm- 1 (scaled for sample thickness.d,

and density, p) and total H20 content of volcanic glass .

Relationship between the band intensity ratio for themolecular-to-total water [Abs(l635 cm- 1)/Abs(3550 cm- 1)]

and total H20 (wt.%) ..

ix

2

4

11

19

23

26

27

31

34

35

36

39

46

50

xLIST OF FIGURES (continued)

14B Relationship between the infrared band intensity ratio for themolecular- to-total water [Abs(1635 em-1)/Abs(3550 em-1)]and Si02 (wt.%) ............................................................ 51

14C Relationship between the infrared band intensity ratio for themolecular- to-total water [Abs(1635 cm- 1)/Abs(3550 cm- l )]and [(Na20 + K20 + CaO + MgO)/(Si02 +Al2~ + P205)] ........ 52

14D Relationship between the infrared band intensity ratio for themolecular-to-total water [Abs(1635 em-1)/Abs(3550 cm-1)]and NBOrr .................................................................. 54

15 Infrared spectra of two water rich obsidian glasses from theca. 1340 AD Mono Craters eruption showing the near-infraredbands at -5200 cm- 1(H20) and -4500 cm-1 (X-OR) in the

6000-4000 cm- l region ..................................................... 60

16 Infrared spectra of two water rich obsidian glasses from theca. 1340 AD Mono Craters eruption showing the presenceof dissolved C02 in the glasses in the 2500-2200 cm- 1 region ....... 62

17 Infrared spectra for Na20-Si02 and K20-Si02 synthetic glasses in

the near-ir region (6000-4000 cm- 1)...................................... 66

18 Infrared spectrum of a hydrated Li2a-Si02 synthetic glass

in the region 4000-2000 cm- l ............................................. 67

19 Deconvoluted Raman spectra for water and hydratedNa20- and K2D-Si02 glasses in the 3000-3800 cm- 1 region ......... 68

20 Raman spectra of water and hydrated K20-Si02 synthetic

glass in the 1500-2ooDcm- l region ....................................... 70

21 29Si MAS NMR spectra for hydrated and dehydratedNa20-Si02 glasses ......................................................... 71

22 29Si MAS NMR spectrum for hydrated K20-Si02 glass.............. 72

23 29Si MAS NMR spectra for hydrated and dehydratedLi20-Si02 glasses .......................................................... 74

24

25

26

27

28

29

30

31

32

33

33

33

33

34

35

36

37

xiLIST OF FIGURES (continued)

23Na MAS NMR spectra for hydrated and dehydratedNa20-Si02 glasses......................................................... 75

7Li MAS NMR spectra for hydrated and dehydratedLi2Q-Si02 glasses.......................................................... 76

Composite sketch of the outer portion of a submarinepillow fragment... .. .. . 80

HTMS H20 release profiles for Loihi (LO-4) layers 1 and 2.......... 82

HTMS H20 release profiles for Loihi (LO-4) layers 3 and 4.......... 83

Infrared spectra for Loihi (LO-4) layers 1,2 and 3in the 1200-4000 cm-1 region....................... 85

H20 release profile from HTMS for Mahukona (MA-21) toshow the annealing temperature..... . . .. . .. .. .. .. . . .. . ... .. ... . . . .. . . . . . . . 87

Infrared spectra of Mahukona (MA-21) glass for unheated andsamples annealed for 1hr, 2 hr and 3 hr, respectively ,. 89

Near-infrared spectrum of Troodos (TR86-9) glass, whichcontains 2.30 wt.% H20, in the 4000-6000 cm- 1 region..... 96

(a) (i) Mica structure. Plan of tetrahedral layer (Si,Al)401O with

tetrahedra pointing upwards, and end view of layer lookingalong y axis.................................................. 109

(a) (ii) Plan and elevation of tetrahedral layer with tetrahedrapointing downwards........................................................ 109

(b) Plan of (i) and (ii) superimposed and linked by a layer ofcations........... 109

(c) Elevation of (i) and (ii) superimposed and linked by a planeof octahedrally coordinated cations........................................ 110

The idealized structure of muscovite 112

Infrared transmission spectra in the OH stretching region ofcleavage plates of phlogopite and muscovite... . . .. . . . . . . . . . . . . . . . . . . . .. . 116

Infrared transmission spectrum in the OH stretching regionof a cleavage plate of biotite...... 118

Diagrammatic presentation of three types of light scattering........... 122

38

39

40

41

42

43

44

45

46

47

48

49

50

51

52

53

54

xii

LIST OF FIGURES (continued)

Energy diagram of a molecule showing the origin of Stokesand anti-Stokes Raman scattering......................................... 122

An example ofchange in polarizability 124

The Raman shift for Stokes and anti-Stokes is always symmetrical.. 124

The frequency of the Raman-scattered light is independent of theexcitation wavelength................................. 126

The origin of Stokes and anti-Stokes effects and their intensitydifference..... 127

Micro-Raman apparatus for use in 1350 and 1800 geometries with aCW laser, and a multichannel Raman spectrograph..................... 133

Raman spectrum of a phlogopite sample (BD3088) in the145-800 cm-1 region........................................................ 145

Raman spectrum of a muscovite sample in the 145-800 cm-1region 146

Infrared spectrum of a phlogopite sample (K17) in the1000-4500 cm-1 region....................... 147

Infrared spectrum of a muscovite sample in the1000-4500 cm-1 region.. 148

Raman spectrum of a phlogopite sample (BD3088) in the hydroxylstretching region (3600-3800 cm- 1) 149

Raman spectrum of a muscovite sample in the hydroxylstretching region (3500-3750 cm-1) 150

Infrared spectrum of a biotite sample (RHP-l) in the1000-4500 cm-1 region............ 151

Basic configuration of the diamond anvil cell. . . . .. . .. . . . . .. .. . . . . . . . . . . . 159

Schematic diagram of (A) Mao-Bell (M-B) diamond anvil cell and(B) M-B cell piston-cylinderdetails. 160

Pressure calibration of ruby R1 fluorescence line to -l60Kbar 161

Ray diagram of refracting beam condenser and DAC for infraredmeasurements............................ 162

xiiiLIST OF FIGURES (continued)

55 Infrared spectra to show the effect of pressure upon theposition of the hydroxyl stretching band for muscovite ................ 165

56 Infrared spectra to show the effect of pressure on the positionof the hydroxyl stretching band for a phlogopite sample (K8)......... 167

57 Relationship between pressure and hydroxyl band positionfor a muscovite sample ........................... '" ....................... 169

58 Relationship between pressure and hydroxyl band positionfor a phlogopite (K8) sample .............................................. 171

59 Relationship between FWHM and applied pressure for thehydroxyl stretching band for a muscovte sample........................ 172

60 Relationship between FWHM and applied pressure for thehydroxyl stretching band for a phlogopite (K8) sample ................ 173

61 The effect of temperature upon the infrared spectrum of a

phlogopite sample (BD3088) in the -1400-1900 cm-l region ......... 177

62 The effect of temperature upon the infrared spectrum of a

biotite sample (RHP-l) in the -1400-2200 cm- l region................ 178

63 The effect of temperature upon the infrared spectrum of amuscovite sample in the -1400-2000 cm- l region ...................... 179

64 Relationship between position and temperature of the hydroxylstretching band for muscovite.............................................. 185

65 Mass pyrograms of a phlogopite mica megacryst (FRB483)showing typical release behavior of (a) H20, F, CI and

(b) CH4, CO, CO2.......................................................... 188

66 Mass pyrograms of a muscovite sample showing the releasebehavior of: (a) H20, HF, F and (b) CH4' CO, S02, C02 ........... 190

1PARTl

CHAPTER I

WATER IN SILICA1E MELTS

A. Introduction

Basaltic magma represents 95 % of the lavas extruded upon the Earth's surface and

seafloor. This partial melt from the mantle provides an invaluable source of information

about the composition and mineralogy of the mantle. Its volatile content has important

implications for modeling the petrological origin and evolution of magmas within the

Earth's upper mantle and crust.

Water is the dominant magmatic volatile within magmatic systems. It is widely

accepted that water plays an important part in the evolution of terrestrial igneous systems.

Consequently, it is important to understand both the thermodynamics and the mechanisms

of silicate melt-water interactions. Ever since its first study by Goranson (1931, 1936) an

enormous number of investigations have been directed to this subject.

Dissolved water profoundly affects physical and chemical properties of silicate

melts. Electrical conductivity and melt fluildity as well as cation diffusivity increase

dramatically with increasing water content (Orlava, 1964; Boulos and Kreidl 1972;

Kushiro, 1978; Watson, 1979; Dingwell and Mysen, 1985; Satherly and Smedley 1985).

These observations led to the conclusion that dissolved water most likely results in

significant depolymerization of the melt (e.g., Kushiro, 1978). This has important

implications for models concerned with magma genesis and eruption mechanisms since the

extent of depolymerization directly affects the ability of a melt to separate from its source,

migrate, assimilate country rock, mix, rise and eventually erupt.

Water, if held within a system under pressure, will also affect phase relations. This

is illustrated in Figure I (Hall, 1987) which depicts the effect of variations in water

pressure on the crystallization of plagioclase. One can see that a melt of composition

AnsoAbso at a temperature of noooc would be in equilibrium with crystals of

composition An9oAblO at P(H20) = 5 kilobars. However, at P(H20) = 0 this same melt

would quench completely to crystals of bulk composition AnsoAbso.

2

1600,----------------------.

1500

I~OO

e.....,"Q.

E111I-

700

Ab An

Ab+An(1 aIm I

Ab+An+H20(PH20=5kbl

Figure 1. The system albite-anorthite-water at high and low pressures (after Hall, 1987).

3Figure 2 illustrates the solubility of water in several petrologically relevant melt

compositions. In all cases, solubility increases with pressure. For some melts it has been

observed to decrease with temperature. Although the wt, % of water within natural magmas

is small, it is the mole % which is of more fundamental importance. For example, 4 wt, %

water in a basaltic magma could be equivalent to a mole fraction of between 10 and 15 %

(Hendersen, 1984).

Burham (1979) stated that in order to fully understand the thermodynamic

properties of the melt, an understanding of the mechanisms of solution of water in silicate

melts is essential. Much effort has been directed towards this end. The albite-water system,

(NaAlSi308-H20) has been extensively studied since albite is one of the end-members of

plagioclase, a crystalline mineral commonly found in volcanic rocks. A wealth of infrared,

Raman, and nuclear magnetic resonance studies have established that water dissolves in

albite and other hydrous silicate glasses in two forms, hydroxyl ions and molecular water

(Ernsberger, 1977; Wu, 1980; Bartholomew et. al., 1980; Stolper, 1982a,b; McMillan et

aI., 1983; Acocella et. aI., 1984; Mysen and Virgo, 1986a,b; McMillan and Remmele,

1986; Farnan et. aI., 1987; Newman et. aI., 1986, 1988; Eckert et. aI., 1988; Silver and

Stolper, 1989; Stolper, 1989; Kohn et. aI., 1989; Mysen, 1990; Silver et. al., 1990).

In addition to water, carbon dioxide and sulfur have also been found to be

important species in volcanic magmas (Anderson, 1975). Other volatiles commonly found

in relatively minor amounts include chlorine, fluorine, carbon monoxide and hydrocarbons.

Even less is known about the dissolution mechanisms of these volatiles in silicate melts.

Although there appears to be a general perception within a large part of the

geological community that water solubility is "understood", even a cursory examination of

the literature shows that this cannot be the case (McMillan and Holloway, 1987). In order

to better understand the interaction of water with silicate melts, various models have been

suggested. In the next section some of these models are briefly outlined.

B. Mechanisms for H20 Solubility

Burnham (l975a,b and 1979) and Burnham and Davis (1974) provided one of the

first quantitative models to explain the dissolutionof water in silicate melts.

12

M.: 10~--o£au,o)0 6I-...J

lD 4:J...Joen 2

1000 2000 3000

PRESSURE .Ibars)

4000 5000 6000

3

2

Figure 2. The solubility of water in silicate melts: (I) Basalt at 1100°C; (2) Andesite at

1100°C; (3) Granite at liquidus, i. e., less than 700°C. (after Hall, 1987).

.+;:..

5Originally, it was generally assumed that H20 dissolved in silicate melts principally

by reaction (hydrolysis) with bridging oxygen's (00) of the melt to produce hydroxyl

groups:

H20 (v) + OO(m) = 20H (m) (1)

where, (v)-----vapor phase; and (m)-----melt phase.

The proportionality between the fugacity of water and the square of the mole fraction of

dissolved water at low total water contents is often cited as evidence favoring the above

reaction scheme. Also, the dramatic increases in the fluidity, electrical conductivity, and

cationic diffusion coefficients with increasing dissolved water are consistent with the above

reaction. This model, however, could not reconcile the fact that the molal solubility of H20

in NaAISi 30g melts was greater than in silica melts at a given pressure and temperature

(Wasserburg, 1957). Therefore, it was proposed (Burnham, 1975a) that, in NaAISi30g

melts, where Na+ provides charge balance on the AI04 tetrahedron in each structural unit,

the dissolved H20 also undergoes dissociation by exchange of a proton (H+) for sodium

NaAISi30g + H20 =HAlSi3~(OH)(ONa) (2)

Here, tetrahedrally coordinated Al is now charge balanced by H+ instead of Na+. One Si-

O-Si bond has been ruptured forming one Si-OH and one Si-O- charge balanced by Na+.

This results in a lowering in viscosity by a factor of 105 to 106, and an increase in electrical

conductivity by a factor of 103 to 104, of the anhydrous melt at the same temperature

(Burnham, 1975a). This reaction was favored for melts with less than 50 mol. % water.

Above 50 mole %, the following mechanism was suggested:

HAISi3~(OH)(ONa) + H20 =HAISi306(OH)3(ONa) .....(3)

i.e., all further dissolved water breaks Si-O-Si bonds to form additional hydroxyl groups.

6This reaction is identical to that suggested earlier (see reaction (1)) to explain the dissolution

of water in silicate melts.

This concept of two different mechanisms to explain the dissolution of water was

later suggested to explain the relationship between the observed molar water solubility in

albite melt and water fugacity (Oxtoby and Hamilton, 1978). Eggler and Rosenhauer

(1978) later confirmed this mechanism for diopside (CaMgSi206)'

Burnham (1975a) also proposed the existence of different hydroxyl complexes,

suggesting that the OH groups may be attached to Si4+ and A13+ in tetrahedral

coordination. Both of these reactions assumed complete reaction with H20.

Mysen et. al., (1980) and Mysen and Virgo (1980) using Raman spectroscopy to

study glasses quenched from water-bearing silicate melts detected hydroxyl groups, but

failed to detect molecular water. So, they assumed the complete reaction of water molecules

to hydroxyl groups. Water was suggested to form hydroxides with Si4+ arid probably Na+

or Ca2+ or both.

Despite the apparent success of models based on the assumption of essentially

complete reaction of water dissolved in silicate groups, considerable evidence from infrared

spectroscopy was available that clearly demonstrated the existence of molecular water

(Ernsberger (1977); Bartholomew et, al., (1980); Takata et. al., (1981) and Wu (1980). An

NMR investigation had also demonstrated the coexistence of molecular water and hydroxyl

groups in silicate glasses (Bartholomew and Schreurs, 1980). The non-linear relationship

observed between f(H20) and the square of the mole fraction of total dissolved water at

high water pressures in synthetically prepared diopside glasses was also used to suggest

the presence of molecular H20 (Eggler and Rosenhauer, 1978). Hodges (1974) suggested,

based on his estimate of the partial molar volume of water in water-rich diopside melts at

P(H20) =20kbar, that molecular water may be present in silicate melts formed at high

pressure with large amounts of dissolved water. Wasserburg (1957), using theoretical

considerations, also proposed that water dissolves in silicate melts as both molecular water

and hydroxyl groups.

7Mysen et. al., (1980) most likely did not detect molecular water because of the

relatively low intensity of the Raman spectrum of water and short integration time (used in

that study) for detecting the principal Raman band at 1600 cm-1 (H-O-H bending), that is

diagnostic of the presence of molecular water. In fact, the Raman spectrum of water is

weak and water is normally used as a solvent in Raman spectroscopic experiments with

little or no interference.

McMillan et. al., (1983) reported the first Raman spectroscopic study to prove the

existence of molecular water within water rich albitic glasses. Later, when Mysen and

Virgo, (1986a) obtained Raman spectra for the system Si02-H20 with longer integration

time and a computerized data aquisition system both H20 dissolved as molecular water and

hydroxyl groups were detected.

Spectra of hydrous aluminosilicate melts differ, however, from those of quenched,

hydrous silica melt in that the 970 cm! band, diagnostic of the presence of Si-OH bonds in

the melts, was not observed in the (AI/AI + Si) range (0.125-0.333) under study. A similar

observation was also made previously by Remmele and McMillan (1984).

In addition to H-OH and Si-OH complexes, hydroxyl complexes may also form

with A13+ or with Na+. Hydroxyl attached to 4-coordinated aluminum (Aliv) results in

Raman bands near 800cm-1 (Ryskin, 1974):

2NaAISi308 + H20 =2NaOH + 3Si032- + 3Si02 + 2A13+ (5)

Non-bridging oxygens are formed in both reactions. (illustrated with the Si032­

stoichiometry). In reaction (4) a portion of tetrahedrally coordinated aluminum from the

three-dimensional network is transferred to the AI(OH)3 complex. In this process, a

proportion of Na+, equivalent to that required to charge balance this amount of Al3+ in the

anhydrous melt, becomes a network modifier as charge balance of the A13+ associated with

8

OH is not necessary. One nonbridging oxygen per newly-formed network-modifying Na+

is stabilized as a result. In reaction (5), for each NaOH, one network-modifying A13+ is

generated. Each such network-modifying A13+ stabilizes three nonbridging oxygens.

Thus, depending on which hydroxyl complex is formed, the rate of depolymerization as a

function of dissolved water will differ.

The majority of the studies on the solubility of water in silicate melts were based

upon alkali aluminosilicates, commonly on the join NaAI02-Si02' Athough, the data in

this system is of fundamental importance for the characterization of the role of water in

magmatic liquids, alkali metals in general (K+ and Na+) and Na+ in particular, are not the

principle cations in magmatic liquids (Mysen, 1990).

Except for rhyolitic melts, 2::. 50% of the metal cations in common magmatic liquids

are divalent (Ca2+, Mg2+ and Fe2+). Infact, over 50% of the divalent cations are

represented by Ca2+ (Mysen, 1988). In order to address this problem Mysen (1990)

examined the interactions between H20 and melts in the system CaO-AI20:3-Si02' He

prepared synthetic samples, where the H20 contents ranged from 3 to 10 wt.% with

Al/(AI+Si) = 0-0.333. and obtained their Raman spectra. Results of these studies were

consistent with water dissolved as molecular water and in the form of various OH

complexes, that included Ca2+ and AI3+. The 970 cm- 1 band from Si-OH stretching

(e.g., Stolen and Walrafen, 1976) observed in the spectra of Si02-H20 melts was not well

resolved in aluminous samples. However, the spectral topology of the fundamental OH

stretch bands near 3600 cm- l could only be rationalized if some Si-OH or (Si,AI)-OH

bonding existed in the melts. Even then, these results indicate that OH groups bonded to

Si4+ is not of major importance in aluminosilicate melts. A similar conclusion was reached

previously by McMillan and Remmele (1986) and Mysen and Virgo (l986a) for melts in

the system NaAI02-Si02-H20.

9Most previous models of water solubility in silicate melts generally assume

essentially complete reaction of water molecules to hydroxyl groups. Evidence for this

hypothesis:

(1) The proportionality between the fugacity of water and the square of the mole fraction of

dissolved water at low total dissolved water contents has often been cited as evidence

favoring models involving essentially complete reaction of water molecules to hydroxyl

groups on solution in the melt. (e.g., Tomlinson, 1956; Kurkjian and Russell, 1958;

Burnham and Davis, 1974; Bedford, 1975).

(2) The dramatic increase in fluidity (decrease in viscosity), electrical conductivity, and

cation diffusion coefficients of granitic melts with increasing dissolved water content are

consistent with solution of water in highly polymerized melts by interaction of water

molecules with bridging oxygens in the melt framework to form OR groups and a less

polymerized melt stucture. (e.g., Burnham, 1975a; Watson, 1979, 1981).

(3) Raman spectra of hydrous silicate glasses show vibrations assigned to OR groups but

none attributed to H20 molecules. (Mysen et. al, 1980; Mysen and Virgo, 1980).

(4) The partial molar volume of water in albite melts at low pressures is similar to the molar

volume of close packed OR- ions, but differs from that of liquid water, perhaps suggesting

that water enters the melt mostly as OH groups rather than as water molecules under these

conditions. (Burnham and Davis, 1971).

(5) The P-V-T equation of state for hydrous albitic melts and the thermodynamic functions

that are derived from an analysis of this equation of state plus available solubility data are

compatible with a model of water in albitic melts that assumes nearly complete reaction of

water molecules to hydroxyl groups (Burnham and Davis, 1974; Burnham, 1975a,b,

1979).

However, these data were found to be inconsistent with a considerable bodyof data

that demonstrated that water-bearing silicate glasses contain both hydroxyl groups and

water molecules. Raman, infrared and nmr spectroscopies have been used to detect

molecular water and hydroxyl groups in silicate glasses: Ernsberger, 1977; Bartholomew

and Schreurs, 1980; Bartholomew et. aI., 1980; McMillan et. aI., 1980; Mysen et. aI.,

101983; Wu, 1980; Takata et, al., 1981 and Stolper, 1982a. Both infrared and nmr provided

quantitative measurements of the concentrations of molecular water and hydroxyl groups in

glasses (Bartholomew and Schreurs, 1980; Bartholomew et. al., 1980; Stolper, 1982a).

For some samples, comparable results were obtained by both techniques.

As we cannot measure the speciation of water in-situ, Stolper (l982b) made the

assumption that the speciation of water within quenched melts is unchanged from that in the

melt and that the speciation isn't a function of quenching history. Because the proportions

of molecular water and hydroxyl groups in glasses were found to vary systematically with

total dissolved water content but showed no apparent correlation with the quenching

histories of the glasses, the concentrations of these species as observed in glasses thus far

probably reflected those of the melt from which they were quenched (Stolper, 1982b).

Direct and indirect comparisons of the structural properties of anhydrous glasses,

melts, and supercooled melts indicated strong similarities between atomic configurations in

these states (e.g., Riebling, 1968; Seifert et. aI., 1981). The absence of microscopic

bubbles in most specimens and comparisons between infrared spectra taken at room

temperature and at liquid nitrogen temperatures strongly suggested that the molecular water

is structurally bound in the glass (Stolper, 1982a) rather than present within fluid

inclusions as a distinct phase.

Stolper (1982a) used infared spectroscopy to provide quantitative measurements of

the concentrations of molecular water and hydroxyl groups within silicate glasses.

In Figure 3 all samples with > 0.5 weight percent total water contain measurable

amounts of molecular water. At 4 wt, % total water content molecular water and water

dissolved as hydroxyl groups are present in approximately equal proportions. Molecular

water dominates at high total dissolved water contents. The following reaction was

suggested to occur at low water contents:

H20 (v) + 0 0 (m) = 20H (m) (6)

OO----bridging oxygen' (-Si-O-Si-)

(v) ---- vapor phase

(m) ---- melt phase

10. I I ( I

1211109

o

8765432

oe Stolper (1982alOB Bartholomrw rt al. 119801l:JJ,. Acocella et at, 1198418

..II>·u~ 6~..Q,

0N

:J:lit-

i 4

A" I I

1I

: , I r2~ Ll ~/ • • •

WI'" H20 In umple

Figure 3. Summary of published OH/H20 distribution in silicate glasses as a function of

total water concentration (open symbols =molecular H20; closed symbols =structurally­bound OH) (after Mysen and Virgo, 1986).

12However, as the total water concentration increased a portion of the water

molecules were thought to dissolve as molecular water:

HZO (v) = HZO (m) (7)

Both species were identified from specific combination bands in the near-ir absorption

spectra of hydrous glasses, and a thermodynamic model was suggested based upon his

results from quantitative infrared absorption measurements (Stolper, 1982a,).

Aines et. al. (1983) showed that their results could be extended to hydrous melts at

high temperatures and pressures. These experiments provided the first direct observation

of the dissolving species in hydrous aluminosilicate melts, and that the thermodynamic

model of Stolper (198Za) and Silver and Stolper (1985) has been more physically

satisfying than the model proposed by Burnham and co-workers, because Stolper's model

is based upon a wider range of data, and on direct physical measurements of the

hydroxylated species. Stolper (198Za) noted that, as the water solubility at high pressure

would be dominated by solvation of molecular HZO as the dissolution mechanism, there

should be little compositional dependence of water solubility at high pressures. Although

the work of Stolper (198Z a,b) has contributed much to the general nature of water

dissolution in silicate melts it did not account for the observed strong compositional

dependence of water solubility over much of the pressure range. Also, for a variety of

hydrous aluminosilicate melts they consistently obtained the correct total water content

(measured independently) for the glass sample by summing the contributions from the

molecular water signal at 5Z00cm -1 and hydroxyl groups at 4500cm-l. This suggested that

within experimental error, all hydroxyl groups in the glass are associated with either Si or

AI, or both. This is an important result, and suggests that the Na-OH species inferred by

Mysen and Virgo (1986 a.b) is not necessary to explain the solubility data. Stolper

concluded,therefore, that models based upon the assumption that water was nearly

exclusively present as hydroxyl groups within silicate melts (e.g. Burnham, 1975a,b;

Mysen et. al., 1980; Mysen and Virgo, 1980) were fundamentally incorrect.

13With the exception of the Burnham model the water solubility mechanisms

discussed so far are based upon infrared and Raman spectroscopic techniques. Although

these techniques provide indirect structural information, they have been used essentially to

determine the speciation of water and not as sensitive probes. Also, the interpretation of

Raman spectroscopic data is debateable. To obtain a much more complete picture of the

structure of glass, nuclear magnetic resonance should be used since it provides information

regarding the spatial relationship between atoms of different types.

Recently, Kohn et. al., (1989) undertook a multinuclear magnetic resonance study

of the structure of hydrous albitic glasses. The structures of a series of hydrous albite

glasses quenched from their melts at high pressures and temperatures were studied using

29Si, 23Na, 27AI, and 1H nuclear magnetic resonance. These glasses were said to be

similar to the glasses studied by others.

The spectra for 29Si and 27AI, and hence their structural environments, were

similar throughout the range of water concentrations studied (0-67 mol.%). However,

major changes in the sodium environment occured. No previous ~odel is consistent with

these results. The data suggest the existence of the following structural features:

(i) Exchange of H+ for Na+ as a charge-balancing cation.

(ii) Formation of Na(OH) complexes.

(iii) Incorporation of molecular water.

(iv) No octahedrally coordinated aluminium.

(v) No AI-OH or Si-OH groups.

These features can be summarised in terms of the equilibrium:

NaAISi308 + H20 = HAISi30g + Na(OH) (8)

The result of fundamental importance from this model is that no depolymerization of the

aluminosilicate framework is required. This study therefore casts doubt upon the model

suggested by Stolper (1982a,b). Okuno et. al., (1987) also showed by using X-ray radial

14distribution analyses of hydrous albite glass that the hydrous glass had a similar structure to

the dry glass, in that, it still consisted of TO 4 tetrahedra and the aluminium was still in

four-fold coordination. These observations agreed with the nmr study. It is obvious that

further work is necessary to provide a better understanding of the interaction of water with

silicate melts.

C. Mechanisms for CO.., solubility

Although water is always considerably more abundant than carbon dioxide in

petrologically relevant silicate melts, the extent to which C02 may effect the siting and

speciation of HZO is an open question. To address this question I very briefly review our

current understanding of COZ dissolution mechanisms.

Our current understanding of the physical and chemical effects of COZ upon magma

is still in its infancy (Fogel and Rutherford, 1990). Raman, infrared and recently nuclear

magnetic resonance studies have been utilized to study these interactions (Mysen, 1976;

Sharma, 1979; Fine and Stolper, 1985a,b; Stolper et. al., 1987; Stolper and Holloway,

(1988); Dixon et. al., (1988); Fogel and Rutherford, 1990; Mattey, 1991 and Kohn et. aI.,

1991). A combination of synthetic and natural samples have been studied to determine the

solubility mechanisms of C02 in silicate melts. As with water, quantitative analyses using

infrared spectrocopy have been carried out (Fine and Stolper, 1985b, and Fogel and

Rutherford, 1990) . At low pressures C02 has been observed to dissolve as carbonate

anions in basalts, whereas in rhyolitic melts it dissolves in the molecular form. In the

former case the COZ molecules are suggested to react preferentially with a non-bridging

oxygen to form carbonate anions:

C03Z-(m) + OO(m) (9)

This results in the polymerization of the melt. Also, COZ can dissolve as molecular C02 in

the melt:

COz(v) = COZ(m) (10)

This is similar to the dissolution of water molecules within silicate melts. However, in

15

albitic glasses C02 dissolves as both molecular C02 and carbonate groups (C032-). The

C02 molecules were recently suggested to reside within "cages" or "holes" (Kohn et. al.,

1991) so that molecular C02 would only be observed if holes large enough to

accommodate the length of the C02 molecule (4.96A) exist. Since, the intertetrahedral

angle for Si-O-Allinkages are probably smaller than Si-O-Si linkages, then the size of the

cage would be expected to increase as the Si/Al ratio increased (Kohn et. al., 1991). Such

considerations could explain why molecular C02 is observed in albite glasses. Other

factors, such as the availability of sites for carbonate formation, could also be important.

The ratio of COl- to molecular C02 may also be a function of the competition between

C032- and A13+ for charge balancing cations such as Na+ (Brearley and Montana, 1~89).

Currently there is no consensus as to the effects of dissolved C02 and H20 on their mutual

solubilities and dissolution mechanisms. Because H20 is always more abundant in

naturally occurring silicate melts it is more likely that the effects of H20 on C02 solubility

will be more important than the effects of C02 on H20 solubility in silicatemelts.

D. The Mass Spectrometric I FfIR Study

Over the past 15 years the High Temperature Mass Spectrometry (HTMS)

laboratory at the University of Hawaii has degassed -1,000 volcanic glasses (mostly

submarine) which range in composition from depleted mid-ocean ridge basalts (MORB) to

highly evolved and enriched andesites and rhyodacites. With HTMS one can obtain mass

pyrograms which permit us to obtain abundances for the volatiles H20, C02, S, Cl, F,

CH4' and CO, among others. These data have allowed us to help constrain models for

subseafloor magmatic processes. For example, glasses from lavas erupted in back-arc

basins are water rich compared to those from mid-ocean spreading centers suggesting a

contribution from the nearby subducted hydrous slab.

From these de-gassing experiments we also obtain information other than that of

strictly petrogenetic importance. Mass pyrograms also give us a "picture" of the degassing

behavior of silicates and hence have the potential to provide information of a more

fundamental nature at the molecular level.

16We would like to interpret these mass pyrograms in terms of volatile

siting/speciation and the physical processes the glass is undergoing as it is being heated. In

this study a combination of high-temperature mass spectrometry (HTMS) and Ff-infrared

spectroscopy (FTIR) have been used to determine the ratio of molecular-to-total water in a

wide range of submarine volcanic glasses. A wide compositional range was chosen to see

what effect it had upon the speciation of water within these submarine volcanic glasses.

The 54 glasses reported here are from mid-ocean ridges and their axial seamounts, back-arc

basins, arcs, and intra-plate oceanic islands.

This is the first study ever to determine the extinction coefficient, for the O-H

stretch band (£3550 em-1), for submarine volcanic glasses using a combination of FfIR

and high temperature mass spectrometry. Previously, a reconaissance study undertaken by

Stolper (l982a) revealed that if the water contents were determined by several different

techniques the extinction coefficients varied considerably. This is due to-interlaboratory

differences and the concomitant differences in the reliability of reported water abundances.

Here, mass spectrometric data were used for all the glasses investigated in this work.

Also, only one publication to date reports a value for the extinction coefficient for volcanic

glasses (Dixon et al., 1988). The method utilized to obtain this value, however, is

unpublished. Further, as already covered in detail in the introduction, there is still

disagreement regarding the dissolution mechanism of water within silicate glasses. This

study is anticipated to further our understanding of this problem.

Only when a series of standards all analyzed by a single technique, can the

limitations and potential of a technique such as infrared spectroscopy be evaluated. Here,

mass spectrometry provided the total water content and infrared spectroscopy provided

information regarding the speciation of water as well as relative water concentrations.

Although, infrared spectroscopy has been used extensively by glass and ceramic

scientists, the same is not true for earth scientists. Hydrogen in minerals most frequently

occurs bonded to oxygen. The resulting OH group is highly polar, making it an efficient

absorber of light in the infrared region thereby making infrared spectroscopy a powerful

tool for the study of hydrogen in minerals and glasses.

17CHAPTERII

EXPERIMENTAL METHODS

A. High temperature mass spectromeny

Quantitative analysis by high temperature mass spectrometry requires the

production of a molecularbeam which is representative of the gases being releasedfrom the

sample. Killingleyand Muenow (1975) state that "such a beam may be produced when the

vapour under study is allowed to emerge from a refractory vessel containing a small

aperture or orifice (Knudsencell)". The sizeof theorifice is critical for studies whicheither

require "free vaporization" or those for which equilibrium be maintained between the

condensed phase(s) and the vapour in the cell (Grimley, 1967; Drowart & Goldfinger,

1967).

The theory behind quantitative analyses begins with an equation for the number of

molecules of species i (Ni) released from a sample maintained. at high temperatures for a

time interval~t (= t2 - t1) with isotope-corrected ion intensity I+i:

t2

Ni = Ki J I+i q dt. (11)t1

where Ci is the average thermal molecularvelocity for species i within the Knudsencell and

Ki is the instrumental sensitivity constant obtained by vaporizing a standard whose vapor

pressure is well establishedover a large temperature range (Killingley and Muenow, 1975;

and Grimley, 1967).

The method used here, to determine the weight percent loss of each volatile, is

independent of an internal standard. However, it must account for all sources of weight

loss from the sample. An example of the calculation is given by Graham (1978). The

relevant equationis given below:

ki Ai Wt. loss of sample--- x x 100% (12)

LKi Ai Wt. of sample

18

~ is the integrated ion intensity area under the (It Tl/2) vs T curve obtained from mass

pyrograms and corrected for fragmentation patterns and isotopic abundances for each

species. Ki is a proportionality constant for each species and involves terms for ionization

cross-section, electronmultiplier effeciency, appearance potentials etc. (Graham, 1978)

The weight loss is determined by weighing the sample before and after the experiment.

Details of this facility (Figure 4) and analytical procedures have been described previously

(Killingley and Muenow, 1975; Liu and Muenow, 1978; Byers, 1984 and Aggrey, 1989).

In a typical degassing experiment, a 40-65 mg sample is placed in a high-purity

alumina crucible, weighed, and heated at a rate of approximately 50C/min. to 12500C

under a vacuum of 10-7 to 10-8 torr. The mass spectrometer monitors degassing of the

sample by rapidly scanning the spectrum from 2 to 100 atomic mass units in 30-second

intervals over the entire temperature range. Except for the possible presence of

submicroscopic vesicles and microcracks which may contain minute amounts of H20 and

other volatiles, the abundances we measure are essentially due to dissolved volatiles in the

glass (Delaney et. aI., 1978; Fine and Stolper, 1985b).

The mass peak due to H20 (rnle = 18) and its corresponding ion-current signals are

measured against cell temperature during each scan; these data are stored on a hard disk for

later data reduction and analysis. When the degassed sample is cooled it is removed from

the mass spectrometer and reweighed to determine total weight loss. System backgrounds

are determined from blank runs; a minimum signal corresponding to a concentration of 5

ppm for H20 is discernible from background contributions. Computer plots of ion­

intensity versus temperature (mass pyrograms. see results section) are generated from the

reduced data. These pyrograms, when corrected for background, permit the quantitative

characterization of sample volatility, since the area under each curve is proportional to the

total amount of H20 released within thecorresponding time-temperature interval.

B. Infrared spectroscopy

Infrared spectroscopy deals with the interaction of infrared radiation with matter.

The infrared spectrum of a compound can provide important information about its chemical

Viewing Port

Forepump

MolecularSieve Trap

TitaniumSublimation Pump

-Vacuum Chamber

Water Cooling--D=E InletShutterplate

--D.::: Knudsen Cell Unit

~ IL II Chamber Inlet (Bleed)Valve

forepump Isolation• Valve

I tdt ~ter Coolingj • Outlet

Gasket

Glove Bag

Ion Pump

Ion PumpIsolation Port

Quadrupole Massfilter

To Quad Control Unit andComputer Inter face

Figure 4. Schematic of high temperature quadrupole mass spectrometer facility. (afterGraham, 1978).

......\0

20nature and molecular structure. Most commonly, the spectrum is obtained by measuring the

transmission of IR radiation, although infrared emission and reflection can also be

measured (Perkins, 1987b).

The infrared region of the electromagnetic spectrum is generally considered to lie in

the wavenumber range from 12,900 to 10 cm- l. Radiant energies in the IR region are on

the order of the energies of vibrational transitions. Hence, IR spectroscopy is one branch

of vibrational spectroscopy. The IR region is often further subdivided into three

subregions. The near-infrared region (nearest to the visible) extends from 12,900 to 4,000

em-I, the mid-infrared region from 4,000 to 400 cm-1, and the far-infrared region from

400 to 10 cm- I.

Nearly all molecules absorb infrared radiation, the exception being homonuclear

diatomics such as 02, N2 and H2' The spectrum is unique for each compound and is,

therefore, useful in compound identification.

Infrared spectroscopy is also useful for quantitative analyses. The absorbance A is

related to the path length d of the sample and the concentration c of absorbing molecules by

the Beer-Lambert law:

.•A:: E cd (13)

where E is called the molar absorption coefficient.

1. Molecular explanation

If two particles of charge +e and -e are separated by a distance r, the permanent

electric dipole moment u, is given by:

I..l = e r (14)

Therefore, heteronuclear diatomic molecules must have a permanent dipole moment since

one atom will be more electronegative than the other and will have a net negative charge.

Homonuclear diatomic molecules cannot have a permanent dipole moment since both nuclei

attract the electrons equally. As a rule of thumb polyatomic molecules with a centre of

21inversion will not have a coincidence between Raman and ir modes. For instance modes

which are i.r; active are Raman inactive and vice versa.

In order to absorb infrared radiation, a molecule must undergo a net change in

permanent dipole moment as a consequence of its vibrational or rotational motion. Only

under these circumstances can the alternating electric field of the radiation interact with the

molecule and cause changes in the amplitude of one of its motions. Quantum theory dictates

that the only transitions that can take place are those for which the vibrational quantum

number, u, changes by unity; that is, the so-called selection rule states that !!J.v = + 1. This

is strictly true for a harmonic oscillator, which has equally spaced vibrational levels.

However, for real molecules we are dealing with anharmonic oscillators. Selection rules

break down under these conditions, resulting in deviations of two kinds. At higher

quantum numbers, !!J.E becomes smaller (distance between levels decreases) and the

selection rule is not rigorously followed; as a result, transitions of !!J.v ± 2 or + 3 are

observed. Such transitions are responsible for the appearance of overtone bands at

frequencies two or three times that of the fundamental band. The intensity of overtone

absorption is frequently low, and the peaks may not be observed.

Vibrational spectra are further complicated by the fact that two different vibrations in

a molecule can interact to give absorption peaks with frequencies that are approximately the

sums or differences of their fundamental frequencies. Again, the intensities of these

combination and difference bands are generally low.

2. Vibrational modes

The relative positions of atoms in a molecule are not exactly fixed but instead

fluctuate continuously as a consequence of a multitude of different types of vibrations.

Vibrations fall into the basic categories of stretching and bending. A stretching vibration

involves a continuous change in the interatomic distance along the axis of the bond between

two atoms. Bending vibrations are characterized by a change in the angle between two

bonds.

The number of possible vibrations in a polyatomic molecule can be calculated as

22follows. Three coordinates (or degrees of freedom) are required to locate a point in space;

to fix N points requires a set of three coordinates for each for a total of 3N degrees of

freedom. Translational and rotational each require 3 coordinates, using up a total of 6

degrees of freedom. The remaining (3N-6) are therefore vibrational modes. For a linear

molecule rotation about the bond axis is not viable, and so two degrees of freedom suffice

to describe the rotational motion. Thus, the number of vibrations for a linear molecule is

given by (3N-5), whereas for a nonlinear molecule there are (3N-6). Each of the (3N-6) or

(3N-5) vibrations is called a normal mode.

Water has the point group C2v and has two hydrogens and one oxygen in its

molecular formula. It is therefore expected to have 3 vibrational modes. Figure 5 shows

schematically the 3 normal modes of vibration for H20. These three vibrations are labelled

vI, v2' and v3. They are the symmetric stretching, bending, and antisymmetric stretching,

respectively of the H20 molecule. The most intense absorption bands of H20 are these

fundamental models. The combination and overtone bands are observed, however, with

lower intensites.

Carbon dioxide is a linear molecule and thus has four normal modes (3(3) - 5) for

vibration. The symmetric stretch in CO2 (VI) does not give rise to a change in dipole

moment and thus is infrared- inactive but active in the Raman spectrum. The other three

modes are infrared-active, but the two bending modes are degenerate and absorb at the

same frequency. Thus C02 shows only two fundamental absorption bands in the IR

region, one due to the antisymmetric stretch (2350 em-I) and the other to the degenerate

bending modes (666cm- I).

3. Instrumentation

In general, a spectroscopic experiment requires a light source, a means of providing

energy resolution of the light before or after interaction with the sample, and a detection

system. Dispersive (double beam) instruments which were used extensively before the

advent of Fourier Transform infrared spectroscopy (FTIR) are rapidly losing popularity in

comparison to FT-IR spectrometers due to the decreasing cost ofFTIR instruments coupled

'\)1 SYMMETRIC STRETCH

'\)3 ASYMMETRIC STRETCH

Figure 5. Modes of vibration of the water molecule

23

24with the inherent disadvantages of dispersive spectrometers. The disadvantages are:

(1) No internal reference exists for frequency calibration, hence instrument requires

external standards for calibration.

(2) Only a small proportion of the energy is available at the detector at any time due to the

monochromator slit, and this may thermally degrade with time.

(3) Long time required for single scan (-30 mins).

(4) High resolution is limited by gratings and slits and the resolution is generally gained at

the expense of spectral intensity.

(5) Stray light can interfere with the measurements.

Fourier transform infrared spectrometers are based upon the principle of the Michelson

interferometer (Perkins, 1986). There are several very signifant advantages that an FTIR

spectrometer has when compared to a dispersive instrument. Perhaps the most significant

of these is the much higher signal-to-noise ratio of the FTIR system (Perkins, 1987a).

Peter Fellgatt (see Perkins, 1987a) pointed out that in a dispersive instrument the spectrum

is examined one resolution element at a time. In contrast, the interferometer sees all of the

wavelengths present all of the time during the scan. The signal-to-noise advantage

(Fellgett's advantage) is equal to the square root of the number of resolution elements in the

wavenumber range that is being scanned:

Fellgett'5 advantage =jv1- v2 (15)

~V

where VI and v2 are the limits of the region scanned and ~v is the resolution. For

example, if we scan from 400 to 4000 cm- l with a resolution of 4 em-I, the advantage

would be a factor of 30. This is only realized when detector noise remains constant as the

signal increases (detector-noise-limited). This is true for thermal detectors such as the

deuterated triglycine sulfate (DTGS) detectors commonly used in most FTIR instruments.

Another advantage is the greater energy throughput compared to dispersive

instruments using a monochromator. This benefit was pointed out by Pierre Jacquinot (see

Perkins, 1987b) and is usually referred to as the Jacquinot advantage. This is obviously of

great use when observation times are limited.

25The other major advantage is encountered with "hard-to-do" samples. Typically

these include microsamples, samples with very low overall transmission, samples

producing very weak spectra, and samples that require the use of low-efficiency

accessories. This is where the interferometer enjoys its greatest practical advantage. By

signal averaging, the noise level is reduced in proportion to the square root of the number

of scans averaged.

C. Sample preparation

1. Mass spectromeny

The samples from which glasses were obtained were from lavas recovered in water

depths ranging from approximately 2000 to 5000 m. These samples span a wide

'compositional range from depleted mid-ocean ridge basalts to water-enriched, highly

evolved rhyodacites. The samples studied in this work were all obtained from their glassey

rinds. Since the lavas were all extruded upon the ocean floor where the mean temperature

is -40 C and erupted at temperatures from between 1100 to l1500C their quenching rates

are believed to be very similar. Additionally this rapid quenching under high hydrostatic

pressure prevents any appreciable loss of volatiles. (see Figure 6). Further, the glassy rind

that fOnTIS on the outside is highly resistent to low-temperature alteration so only minor

sample pretreatment is necessary before analyses. Only the best glass is chosen to ensure

that the glass is as close to the chemistry of the magma that existed before its extrusion.

Glasses free of palagonization, MnO crusts, phenocrysts and vesicles were used in

this study. Glass from the outermost rind of lavas was chiseled off and crushed in an agate

mortar and sieved between 16 and 32 mesh (0.5-1.0 mm in diameter). The freshest

selected glass fragments (approx. 80 per run) were washed in 0.3 N HCI to remove any

surface adsorbed carbonate and ultrasonically cleaned in doubly distilled water. The cleaned

samples were dried in an oven at 1200C for 12 hours and cooled in a desiccator. Dark­

brown glassy fragments free of phenocrysts and vesicles were handpicked using a

binocular microscope and then degassed in an effusion-vaporization source (Knudsen cell)

interfaced with a computer monitored quadrupole mass spectrometer. The need to clean

and carefully pick the samples is clearly justified by looking at Figure 7a,b, which shows

26

E.Jo:

0 0

o

(a)

o

o

4 o

o

~ 0

o 06

Wt % ~O

Figure 6. Weightpercent 1120 as a function of ocean depth (after Killingley and Muenow,1975).

27

--~--- ----,------..,....------.,

600 BOO 1000

(a)

1II

~~

~0~~1~ I

1I

1I

--~1200

Tcmpel"atul"e (el

-450N East Pacific Rise Ridge Crest

9. 6~---~---r- ----,-------- ..-,--.. -----.

9.S

"to.; 9,4

c41...: 9.3

CQ.

9.2

Muhukona (MA-21)

----~

(b)

12\lt0600 0'''"Tel'lperat~lrp. (el

'l.lilk.......---------------=-----'-'---..-.,;Figure 7. Mass pyrograms showing the release profile of water for (a) altered glass and (b)unaltered glass.

28the difference between a carefully picked versus a sample picked without due care. The

low-temperature H20 peak centered at about 4500C (Figure 7a) is probably due to

degassing from a clay mineral (e.g., Kaolinite AI2Si20:3(OH)2), or a hydrate such as

gypsum (CaS04.2H20).

2. Infrared

Glass grains (0.5-1.0 mm in diameter) prepared for use in the high temperature

mass spectrometric analyses were also used for all infrared analyses. Approximately 2 to 3

grains of each sample were mounted within a specially made phenol based ring (outer

diameter 12 mm, inner diameter 1Omm, thickness 2 mm) using orthodontic resin (L. D.

Caulk Co.). The ring allowed for easier sample handling. Balsam adhesive (balsam

neutral, Buehler; No. 40-8110-004) was used to affix these rings onto microscopic slides.

One side was ground down to - 300 urn using an Ingram thin-section grinder. Final

polishing was accomplished using 0.3-1.0 urn diamond paste on a lapping wheel. Once

one side was polished satisfactorily, the ring was flipped over. To ensure the polished side

was firmly pressed against the microscopic slide face the whole slide was left underneath a

press while the adhesive set. The other side was ground and polished similarly. Polished

grains were removed from the binding resin by cleaning with acetone in an ultrasonic bath.

They were stored in air prior to infrared analyses. Sample thickness was measured using a

digital indicator (543 series Digimatic Indicator, Mitutoyo Manufacturing Co., Ltd) with a

resolution of 2-3 urn.

D. Instrumental Methods

A Perkin-Elmer 1720X Fourier Transform Infrared spectrometer was used to obtain

transmission spectra in the 2.5- 25 urn (4000-400 cm-1) wavelength region. In order to

avoid visible, vesicles or phenocrysts the sample was placed over a -600 urn aperture. The

aperture was placed within a beam condenser accessory. All spectra were obtained with a

resolution of 4cm-1 and scan speed of 0.5 em/sec using a DTGS detector. Generally 300

29scans were found to be adequate for each spot studied. All spectra were converted into

absorbance mode for quantitative analyses.

E. Density measurement

Glass densities were measured on bubble-free glass fragments using either the

float-sink method or with a pycnometer. For most of the measurements the flotation

method was utilized. This method involved using two halogenated organic heavy liquids

with the densities 2.00 g/ml and 2.95 g/rnl (tolerances of ± 0.005 g/ml), from R. P.

Cargille laboratories, Inc., New Jersey.

The flotation method entails using two miscible liquids in which the sample grain is

insoluble (one more dense, one less dense than the sample). The two liquids are mixed in

such proportions that the grain remains suspended; that is, it neither sinks nor rises to the

surface of the resulting mixture. The density of the sample is then obtained by weighing 1

ml of the liquid mixture, that contained the suspended samplegrain..

The pycnometer is a small bottle fitted with a ground glass stopper through which a

capillary opening has been drilled. In order to obtain the specific gravity the following

measurements are required:

Wl-----weight of empty bottle

W2-----weight of pycnometer + sample

W3-----weight of pycnometer + sample + distilled water

W4-----weight of pycnometer + distilled water

(W2 - WI )----- weight of sample

W4 + (W2 - WI) - W3-----weight of water displaced by the sample

Specific gravity = (W2 - W1) (16)W4 + (W2 - WI) - W3

To obtain W4, the pycnometer is initially only partially filled with distilled water and boiled

for a few minutes to drive off any air bubbles. After cooling the pycnometer is funher

filled with distilled water and weighed. When filling with water it is important that the

water rises to the top of the capillary opening but that no excess water is present. The

density of the sample is obtained by multiplying the specific gravity by the density of

30distilled water at the same conditions at which the specific gravity was measured. Both

techniques gave results that did not vary by more than 2% ( ± 50 gIL).

F. Magic angle spinning nuclear magnetic resonance (MAS NMR)

Nuclear magnetic resonance spectroscpy (NMR) is a powerful probe of the static

structure and dynamic behavior of condensed phases. It has been used for many years by

chemists as a standard tool for the determination of the structures of organic molecules in

solution, by obtaining their 1Hand l3e nmr spectra. However, in the past the nmr spectra

of solids in general gave broad bands, this made nmr of limited use for studying solids.

With the advent of high-field superconducting magnets, magic-angle spinning (MAS), and

the ability to examine nuclides with a quadrupole moment has made routine study of a wide

range of solids possible.

The broadness of the bands for solids is attributed to a combination of interaction

terms. These terms contain (3cos 26-1), where 6 is the angle between H 0 (the applied

magnetic field) and the principle axis of the interaction. When the sample is spun at 9 =

54.70 most of these interaction terms go to zero, resulting in line narrowing. This

narrowing due to MAS occurs whether the sample is crystalline or amorphous. However,

the bands for amorphous samples are not as narrow, because of static structural disorder

present in the amorphous and glassy samples.

29Si is the most thoroughly investigated and widely applied nuclide of .

mineralogical interest. Its study was previously restricted because of the low sensitivity of

the 29Si nrnr method. With the introduction of pulse Fourier transform nmr has opened the

way for its routine study. The 29Si NMR spectra give characteristic and mostly well

separated signals for the Si04 groups in different structural environments. For instance

nmr spectroscopists use the Qn notation to denote the connectivity of the number of other Q

units attached to the Si04 tetrahedron under study. Thus, Qo denotes the monomeric

orthosilicate anion SiO44-, Q 1 end-groups of chains, Q2 middle groups in chains or

cycles, Q3 chain branching sites and Q4 three-dimensionally cross-linked groups (see

(°0)

O-J

- o-Si-Q-I0-

(°2)

0-I

Si-G-Si-O-SiI

0-

(°1 )

0-I

-O-SK)-SiI0-

(°3)

Si

~S i-G-9t-Q-S i

0-

31

Figure 8. MAS NMR Qn notation for silicate anions. (where n indicates the connectivity,i.e., the number of other Q units attached to the Si04 tetrahedron).

32Figure 8). From their signal intensities the relative concentrations of the different structural

entities can beestimated.

The MAS NMR measurements on hydrated alkali silicate glasses were carried out in

collaboration with Dr. Barbara L. Sherriff at the Prairie National Centre, Winnepeg,

Canada. All spectra were obtained from powdered samples on a Bruker AMX-500

multinuclear Fourier transform nmr spectrometer which has an additional magnet of 360

MHz. 29Si, 7Li and 23Na spectra were recorded at 99.4 MHz and 132.3 MHz

frequencies, respectively. The sample was spun at an angle of 54.70 to the magnetic field

using an MAS probe of the type described by Fyfe et. al., (1982). Several specially

prepared synthetic hydrated glasses and samples of the same glasses dehydrated by

remelting were studied by this technique.

33CHAPTER III

A. RESULTS AND DISCUSSIONA. Results

A typical mass pyrogram (ion-intensity versus temperature) for the release of water

(rn/e = 18) from a submarine volcanic glass is shown in Figure 9. This particularpyrogram

is from a glass taken from the outermost rind of a pillow basalt from Loihi seamount,

Hawaii dredged from a depth of approximately 2 km. This pyrogram is generated by a

IBM-PC/AT computer after data reduction. When corrected for background, it permits the

quantitative characterization of sample volatility since the area under each curve is

proportional to the total amount of the volatile released within the corresponding

temperature interval. Figure 10 shows the pyrograms for the release of H20, C02 and

S02 for a submarine glass sample from the Juan de Fuca Ridge, located off the northwest

coast of the state of Washington.

Analytical uncertainty for H20 based on reproducibility is + 0.020 wt %. Accuracy

is difficult to judge due to the lack of suitable standards. However, analyses made on splits

of a representative basaltic glass (glass P1505, Table 1) from four different laboratories

using different experimental techniques (Peter Michael, private communication) gave the

following results: 0.51 ± 0.025 wt %, CHN gas chromatography; 0.41 ± 0.02 wt %,

manometric H2 gas; 0.50 ± 0.02 wt %, FTIR spectroscopy; 0.455 ±0.020 wt %, mass

spectroscopy, this laboratory.

A typical infrared spectrum of a submarine volcanic glass is shown in Figure 11.

The two bands of interest in this region appear at -1635 and 3550 em-I. The band at 1635

em-I is the fundamental bending mode of molecular water; the broad asymmetric band at

3550 cm-1 is due to the fundamental stretching of hydrogen with respect to oxygen and

therefore has contributions from both molecular water and other OH-containing species

(Nakamoto, 1986). The breadth and asymmetry of the band at 3550 cm- 1 reflects the

envelope of contributions both from the symmetric and antisymmetric stretching modes of

hydrogen in H20 molecules. The broadening of the band also indicates additional

1.00. iii iii iii iii i i

H20+

0.80

0.60

0.40

0.20

1250850

TEMP (OC)

4500' , Ills: !! ,!,! I

Figure 9. Mass pyrogram for a Mahukona (MA-21) submarine basaltic glass showing [hedegassing behavior of H20 (m/e = 18) with respect to temperature.

Yo)

~

1200

.600

.500II1'~ F-- H20

~ \"

., 40 I-rI.O /'~ .c '~ I~ I

H .300 Ic "~ I

.200 I- JI '\I., 1",)'1J I f rt, J J1lW

'. if II ~trJ.,r,,-· , '

• 100 h-\,l'lMhjJJi""\vtl1,J'!'t\"iv,''''\'/MIIf\'I.;\f~(V~\.;V''~ll\/f\J,~!. _,J)

SO.,-'\.,

400 600 800~----:~':=:=----:~~-.J

Temperature eelFigure 10. Mass pyrograms for a Juan de Fuca (JDF-21) basaltic glass showing thedegassing behavior of H20, C02, and S02 with temperature.

wVI

36

2.0

wuz«en0::o(f)

en« 1.0

03 80 0 2600 1400

WAVENUMBERS (ern")

Figure II. Infrared absorption spectrum of a typical submarine volcanic glass between1400-4000 cl11- 1.

37contributions from the distribution of hydrogen-bond strengths of the hydroxyl and H20

groups dissolved in the glasses (Scholze, 1959; Paterson, 1982).

Combination and overtone bands in the near-infrared region have been utilized

previously (Silver and Stolper, 1989, Newman et. aI., 1986) to calculate extinction

coefficients for the fundamental bands in hydrous albitic and rhyolitic glasses. However,

these bands were not observed for glasses in the present study due to the lower water

contents and the presence of strong charge-transfer bands caused by the presence of iron

ions in the glass. Two of the bands observed in this region for some hydrous

aluminosilicate glasses would have been of particular interest. One of these occurs near

4500 cm- I which arises from the combination of O-H stretching vibrations near 3550

em-land Si-OH (AI-OH) vibrations between 900 and 1000 cm! (Stolper, 1982a). The

4500 cm- I band is therefore due solely to hydroxyl groups. A second band at -5200 cm-1

is due to a combination of the H20 bending vibration near 1635 cm-1 and the O-H

stretch of molecular water, and is therefore due only to molecular water. The inability to

observe these two near-infrared bands in the spectra of natural glasses precludes one from

obtaining seperate values for abundances of molecular water and all other OR-containing

species. Therefore, the molecular water-to-total water ratio is reported here. A set of

water-rich glass samples from the Troodos ophiolite (Cyprus) where these bands were

observed well enough to be able to calculate the extinction coefficients for these ~wo bands

in the near-ir region will be discussed later.

Not observed in Figure II are the carbonate doublet peaks generally centered at

approximately 1530 and 1435em-I for C02-containing basalticglasses (Fine and Stolper,

1985b). Even though the sample whose spectrum is shown in Figure 10 contains 700 ppm

C02 (Byers et. aI., 1983) these were not observed due to the very thin sample sections

required to accurately measure the 355Ocm- 1 band in this water rich glass. In lower water

containing glasses where the 3550 em-I is less intense and did not completely absorb the

infrared radiation it was possible to use thicker samples (see Table 2). For these glasses the

38carbonate doublet band was clearly observed. In none of the basaltic glasses studied was

a band at 2350 cm- 1 attributable to the presence of molecular C02 observed. By careful

examination of the infrared spectra of all the samples studied it appeared that for those

samples that were rich in molecular water the carbonate bands were not observed, even

when the C02 concentration was high. However, when the molecular water content was

low the carbonate bands were clearly observed. Figures 12 clearly depicts this observation.

Table 1 gives the major elements, water abundances and sample locations for all the

glasses included in this study.All the relevent experimental data necessary to calculate the

extinction coefficients for the fundamental OH-stretch at 3550 cm- 1 and the peak heights

for the molecular water bending mode (1635 em-1) are shown in Tables 2 and 3,

respectively. Table 4 gives the major element and infrared spectroscopic data for two silica

rich glasses.

Band intensities at 1635 cm- l and 3550 cm- 1 were measured using the base-line

method (Hawthorne and Waychunas, 1988) This involves drawing a straight line between

the minima for the absorbance band of interest; the distance between the maximum and the

point where a vertical line intersects the baseline is then defined as the absorbance of the

band. In cases where aluminosilicate network vibrations interfere with a smooth

interpolation of the baseline, spectral subtraction of a reference sample void of absorbing

species is commonly used. For example, Newman et. al., (1986) in a study of water in

rhyolitic glasses found that a silicate combination/overtone at 1600 cm- 1 interfered with the

1635 cm- 1 molecular water band and therefore employed spectral subtraction using a

vacuum-melted anhydrous glass of similar composition. Since most of my samples are

basalts and andesites these silicate bands are small and interference was found to be

minimal. Spectral subtraction was tried for those few samples (e.g., 994-1D) where these

silicate bands showed significant intensities. However, the background for the vacuum­

degassed samples changed significantly in the region of interest thereby making this

H..,O BE\DIi':G

~

C031- BA:\DS

1fT.:': C02 i~'II

A 0.698 WT.:r: H20 .:\~D 0.-1-33

B 0.898 WT.~~ H20 .\ND 0.268 WT.~ C02

.5~(J

~~

..0So4oen

..0-.:

W\0

3000 2000Wavenurnbers (ern-i)

Figure 12. Infrared spectra for Mariana arc volcanic glasses in the 4000-1200 <:111- 1 region.

a , I I J

-1-000

Table I. Major elements and H20 abundances for volcanic ulasscs (wt, q,}

Sample * H2O SiOz Ti02T MgO CaO Na20 K20 P205 Mg# ReferenceALZ03 FeO

Spreading Centers and Axial SeamountsCDll-12 (1) 0.881 50.06 1.78 IlW2 7.55 6.96 9.65 3.76 1.23 0.52 65 Aggrey ct al., 1988aCD6-B2 (I) 1.021 50.61 2.41 17.58 7.85 4.44 8.32 4.7H 2.16 0.77 53 ibid.CD6-B1 (1) 0.357 50,48 1.75 15.19 9.91 6.86 11.74 3.26 0.22 0.25 58 ibid.CD·2 (1) 0.606 49.44 2.01 16.70 9.21 6.75 10.74 3.10 0.73 0.30 59 ibid.CD12-1 (1) 0.675 48.90 1.69 17.80 8.18 7.69 10.48 3.38 0.76 0.39 65 ibid.CDl-6 (I) 0.126 51.12 1.17 14.52 10.65 7.72 12.59 2.23 0.06 0.12 59 ibid.CDll-5 (I) 0.792 50.00 1.79 17.80 7.78 6.96 9.70 3.69 1.19 0.54 64 ibid.ALV91OR4 (2) 0.100 50.10 1.27 16.18 8.98 8.23 12.23 2.56 0.05 0.10 64 Byers et al., 1986ALV91 IRS (2) 0.190 50.16 1.50 15.44 9.66 7.68 12.00 2.79 0.11 0.13 61 ibid.ALV916R8(2) 0.260 50.28 1.58 15.13 9.75 7.14 11.79 2.90 0.27 0.17 59 ibid.P1505·1 (3) 0.455 51.10 1.56 14.76 10.31 6.95 11.56 2.42 0.50 0.19 57 Michael and Chase, 1987EN-112-4D-6 (4) 0.362 49.80 2.22 14.11 12.03 6.58 11.27 2.84 0.14 0.21 52 ibid.TTI52·11·7 (5) 0.160 47.90 1.30 16.80 10.15 9.41 11.82 2.80 0.03 0.11 63 Dixon et al., 1988TT152-61 (5) 0.270 49.20 1.70 14.60 11.86 6.83 12.03 2.90 0.16 0.16 51 ibid.TT152-21 (5) 0.280 50.60 1.90 14.20 12.00 6.70 11.00 2.60 0.16 0.18 48 ibid.TT152·13 (5) 0.250 50.50 2.40 12.95 14.10 5.90 10.90 2.70 0.15 0.23 43 ibid.TTI52·1I·21 (5) 0.260 50.50 l.80 13.90 11.70 6.92 12.03 2.30 0.21 0.21 52 ibid.HU81017-ll (5) 0.220 51.10 1.50 15.30 9.30 7.50 11.80 3.20 0.10 0.16 59 ibid.HU81017-2-4 (5) 0.370 51.60 I.lJO 14.20 11.11 6.23 11.18 3.00 0.44 0.33 50 ibid.CHNllS-4 (6) 0.680 SO.70 1.30 14.90 7.70 8.10 13.30 2.20 0.40 n.d, 68 Stolper, 1982aD6·C44 (7) 1.170 57.35 1.80 13.30 12.30 2.67 7.14 3.33 0.56 0.18 30 Byers ct al., IlJ84994·3A (7) 0.940 55.90 2.60 11.50 15.60 2.50 7.30 3.20 0.22 0.54 24 Byers et al., 1983994-10 (7) 0.770 53.10 2.70 12.20 14.78 3.90 8.60 2.80 0.19 0.33 34 ibid.999·IB (7) 1.270 56.30 2.00 11.10 16.30 1.60 6.80 3.00 0.25 0.72 16 ibid.loo2-4B (7) 0.870 54.30 2.60 12.00 15.50 3.50 8.20 3.00 0.20 0.47 31 ibid.Back-Arc Basins32·7 (8) 0.638 48.41 2.32 16.95 10.08 6.39 9.68 5.02 0.07 0.32 56 Muenow ct al., 199132-8 (8) 0.590 48.26 2.31 16.97 10.02 6.47 9.78 5.08 0.09 0.32 56 ibid.32-9 (8) 0.631 48.14 2.32 17.12 9.H3 6.36 9.72 5.26 0.07 0.32 56 ibid.32-10 (8) 0.562 48.27 2.21 16.95 9.93 6.45 9.47 5.04 0.07 0.32 56 ibid.26-8 (8) 0.221 51.15 1.64 14.86 10.00 7.49 11.22 2.88 0.08 0.20 60 ibid.26-P (8) 0.261 50.70 1.60 15.68 9.29 7.84 11.76 2.88 0.10 0.24 63 ibid.

~0

Table I. (Continued) Maior cJcmenl~~II)JiH20!lbunclancesJorvolcaniculasscs (WI. %)

Sample * H2O Si02 TiOZ-T--

CaO Na20 K20 P205 MglI ReferenceAI203 FeO MgO

Back-arc Basins (Continued)19-7 (9) 0.800 49.25 2.04 15.85 10.13 6.41 10.93 3.00 0.68 0.21 53 Aggrey ct al., 1988hD20-43 (10) 0.733 50.70 1.29 16.58 7.69 7.67 11.12 3.15 0.3-1 0.19 67 Muenow et al., 1980020-35 (10) 0.717 50.62 1.28 16,49 7.57 7.64 11.05 3.18 0.30 0.18 67 ibid.D23-4 (10) 0,480 50.34 1.45 16.45 8.27 7.32 10.85 3.44 0,43 0.20 64 ibid.GUI-8 (I I) 1.152 50.30 0,40 15.74 7.07 8.25 13.68 1.48 0.08 0.13 70 Aggrey, 1989G145-2 (11) 1.738 64.30 0.87 11:65 11.31 0.50 5.13 3.97 0.29 0.33 8 ibid.MV1349 (12) 0.898 50.05 1.22 16.16 8.91 7.38 11.4 2.88 0.50 0.13 63 Garciaet al., 1979Mariana ArcMVI5220 0.898 56.50 1.01 15.52 9.25 2.98 8.12 2.83 1.07 0.22 39 Garciaet al., 1979MV1514 0.837 55.40 0.96 15.94 9.57 3.57 9.16 2.67 1.06 0.18 43 ibid.MV15247 1.034 55.55 0.98 16.02 10.25 3.33 8.32 2.84 1.16 0.22 40 ibid.HawaiiKK19-20 0.800 47.23 3.34 13.83 11.88 5.95 11.51 3.15 0.90 0,44 50 Byers et al., 1985KK30-1O 0.830 46.51 3,43 13.10 12.59 6.52 11.88 3.30 0.92 0,48 51 ibid.KK17-2 0.840 46.90 3.37 14.07 12.46 5.64 11.67 3.14 0.87 0.39 52 ibid.KK18-8 0.640 48.48 2.84 12.38 11.93 7.33 12.60 2.37 0.51 0.27 55 ibid.MKI-8 0.745 52.49 3.18 13.66 10.87 5.26 9.22 2.87 0.75 0.41 49 Garciaet al., 1989MK5-2 00407 51.71 3.25 13.25 12.31 5.54 9.87 2.52 0.60 0.37 47 ibid.MU-IO 0.350 52.38 2.83 13.58 11.05 6.01 9.86 2049 0.54 0.33 52 ibid.ML2-3 0.175 52.65 2.26 14.10 10.06 6.87 10.88 2.37 0.36 0.27 57 ibid.ML3-30 0.548 52.07 2.32 14.53 10.37 6.30 10.59 2.32 0.37 0.26 55 ibid.ML4-2 0.259 51.90 2.65 13.58 11.41 5.98 10.25 2.36 0.49 0.34 51 ibid.LO-4 0.588 48.69 2.82 13.44 11.82 6.81 11.68 2.54 0.54 0.32 53 ibid.MA-21 0.692 47.72 2.83 15.30 11.86 6.25 11.59 2.81 0.48 0.34 51 Garciact al., 1990MA-32 0.547 47.12 2.76 15.25 11.87 6.71 11.60 2.78 0.48 0.34 53 ibid.

*Numbers in parentheses refer to locations as follows:

(I) CocosSeamounts, East PacilicRise; (2) East PacificRise, 21~ (3) Explorer ridge, 500N; (4) East Pacific Rise,330S;

(5)Juande FucaRidge; (6) Mid-Atlantic Ridge; (7) Galapagos Spreading Center; (8) Woodlark Basin; (9) Fiji/Basin;(10)ScotiaS<:<I; (II) Manus Basin; (12) Mariana Trough. +>.......

Tahle 2. Data for fundamental OH-Slretch peak C355Ocm-l)

Sample H2O Absorbance Thickness Density Abs. Coeff.(wt, %) (Peak height) (urn) (giL) (1/I11ul-cl11)

CD11-12 0.881 1.119 152 2735 55CD6-H2 1.021 1.006 114 2718 57CD6-81 0.357 0.432 155 2783 51CD-2 0.606 0.486 85 2752 62CD12-1 0.675 0.442 70 2787 60COI-6 0.126 (J.078 82 2791 49COIl-5 0.792 1.529 186 2725 69ALV91OR4 0.100 0.112 96 2789 75ALV911R5 0.190 0.130 68 2847 64ALV916R8 0.260 0.622 200 2831 76P1505-1 0.455 0.672 170 2788 56EN-112-4D-6 0.362 0.274 100 2828 48TTI52-77-7 0.160 0.414 288 2833 57TT152-61 0.270 0.819 249 2848 77TT152-21 0.280 0.686 195 2804 81TTI52-13 0.250 0.890 302 2845 75TTI52-11-21 0.260 0.527 193 2777 68HU81017-11 0.220 (1.668 234 2813 83HU81017-2-4 0.370 0.889 223 2840 68CHN115-4 0.520 00402 75 2767 67D6-C44 1.170 0.320 35 2561 55994-3A 0.940 0.644 65 2597 73999-18 1.270 1.565 125 2621 681002-48 0.870 0.376 64 2632 46994-10 0.770 0.279 35 2585 7232-7 0.638 0.726 133 2800 5532-l! 0.590 0.811 148 2769 6032-9 0.631 0.487 88 2820 5632-10 0.562 0.867 157 2766 6426-8 0.221 0.375 169 2X55 6326-P 0.261 0.370 168 2810 5419-7 0.800 1.371 180 2789 62D20-43 0.733 2.000 296 2795 59D20-35 0.717 0.750 108 2803 62023-4 0.480 0.290 84 2796 46GL31-8 0.152 0.746 74 2784 5JGL45-2 1.738 1.516 103 2574 59MV1349 0.898 0.568 62 2556 72MVI5220 0.898 0.385 56 2430 56MV1514 0.837 1.547 252 2535 52MV15247 1.034 2.010 250 2539 55KK19-20 0.800 1.343 174 2893 60KK30-10 0.830 0.977 115 2869 64KK17-2 0.840 0.515 60 2831 65KK18-8 0.640 0.767 120 2888 62MKI-9 0.745 . 0.761 109 2681 63MK5-2 0.407 0.742 194 2732 62MLl-I0 0.350 0.604 202 26111 57ML2-3 0.175 0.121 72 2713 64MU-30 0.548 1.314 223 2815 69MlA-2 0.259 0.571 201 2643 75Lrn 0.588 0.970 152 2965 66MA-21 0.692 1.548 205 2884 68MA-32 0.547 1.170 183 2896 73

42

Table 3. Data for fundamental H20 bending mode peak <1635 cm-1)

Sample H2O Absorbance Thickness Density

(wt. %) (peak height) (11m) (gIL)

GL31-8 1.152 0.063 74 2784GL45-2 1.738 0.508 103 257419-7 0.800 0.076 180 2789MA-21 0.692 0.120 205 2884"MA-32 0.547 0.056 183 2896LO-4 0.588 0.038 152 2965994-3A 0.940 0.122 65 2597999-lB 1.270 0.296 125 2621994-1D 0.770 0.040 35 2485D6-C44 1.170 0.069 110 2561020-35 0.717 0.019 108 2803020-43 0.733 0.105 296 2795KK19-21 0.800 0.064 174 2893KK30-10 0.830 0.044 115 2869KK17-2 0.840 0.026 60 2831KK18-8 0.640 0.044 120 2888MV1349 0.898 0.028 62 2556MV15220 0.898 0.121 56 2480MV1514 0.837 0.311 252 2535MV15247 1.034 0.394 250 2539C06-B2 1.021 0.063 114 2718

43

Table 4. Data for low water-high silica glasses

44

Sample

Si02Ti02

A1203FeOMgOCl0Na20

K20

H20Density (gIL)

Thickness (urn)AbsorbanceAbs. Coeff.(l/mol-cm)

Obsidian

76.930.0512.510.840.020.473.894.59

0.0732436267, 1590.528, 0.321200,205

Tektite

72.830.712.914.521.81.641.252.560.006

24215060.082201

45procedure unusable. For these samples thin sections were made as thin as possible to

minimize this interference.

The Beer-Lambert Law can be used to express the relationship between wt. % of

H20 in a glass and its absorbance:

Wt.% = 18.02 x abs. x 106 (17)pxdxe

By substituting the measured sample thickness (d urn), density (g IL), absorbance (abs.)

and the mass spectrometric water value, it is possible to calculate the molar absorptivity (c)

(or extinction coefficient) for the fundamental OH-stretch at 3550 cm- 1 for each glass

studied. Plotting (abs.)/(p x d) against wt. % H20 (Figure 13) a "best fit" value for e has

been obtained. Based upon uncertainties in total water (± 0.02 wt. %), absorbance C±

0.02), sample thickness (± 2 urn) and density (± 50 gIL) a value of 61±1 L/mol-cm was

obtained. This plot was constrained to go through the origin because for Beers law to be

valid, at zero water content, there should be no absorbance.

The major causes for the scatter shown in Figure 13 were considered. There is no

apparent relation between the molar absorptivity reported in Table 2 and the major element

chemistry, total water content or the tectonic setting from which the glass was obtained.

The specific location of the individual data points in Figure 13 are, however, dependent :

upon the mass spectrometric water abundances but since these were all determined in this

laboratory the variation in molar absorptivities cannot be due to systematic differences in

the measured water abundances. It is possble that the variation in the quenching rate of the

magma on the seafloor is affecting these data but this is difficult to assess. For example,

from a depth profile study (see Appendix C), it has been observed from a few infrared

measurements made on glass taken from between 10 to 20 rom below the outermost glassy

rind of pillow lavas, that there is an apparent increase in the proportion of molecular water

from these more slowly cooled regions of the lava. However, since most submarine

magmas are believed to erupt at similar temperatures (from between 1100 to

46

2.0

o

O~ __L ....l._ ____JL...._ _=_'

o

6.0r------------------~---,

<D 4.00 0

>< 0

-0Lf)io "Uro

Q...........(/)

..0« 2.0

Figure 13. Relationship between the band intensity for the fundamental OH-stretch at-3550 cm- 1 (scaled for sample thickness, d, and density, p) and total H20 content ofvolcanic glass.

47

12000 e for basalts, Rodey Batiza, private communication) and onto the seafloor where the

mean temperature is -40 e worldwide, it could be argued, that at least for glasses sampled

from the outermost rind of quenched lavas, the effects of slight differences in temperature

and cooling histories on water speciation would be small. Heterogeneous distributions of

water in some poor quality glasses may also be a source of additional uncertainty.

However, there is considerable variation in s-values for the OH-stretch reported in the

literature. The E values range from 33 to 186 L/mol-cm depending upon glass composition

(Stolper, 1982a; Shelby et. al., 1982 and Bartholomew, 1983). The higher values (77 to

186 L/mol-cm) of E are generally those associated with vitreous silicas. For basaltic

glasses similar to those studied here Dixon et. al., (1988) obtained a E value of 63 ± 5

L/mol-cm which is within the experimental uncertainty of our value of 61 + 1 L/mol-cm. It

is interesting to note that the stretching band envelope for liquid water has a value of 81

L/mol-cm (Thompson, 1965).

It is conceivable that the variation in s-values for the 3550 cm- 1 band reflects a

competition among tetrahedral (Si,AI) and non-tetrahedral cations (e.g., Na) for hydroxyl

in the melt. For example, if all OH groups are associated with silicon then on vacuum

pyrolysis the glass will produce one mole of water for every two moles of OH, viz. ,

2 (Si-OH) = Si-O-Si + H20 (18)

or, more generally,

20H(melt)= 0 0 + H20 (19)

where 0 0 is a bridging oxygen. Therefore, on a weight basis, 34 pans of OH in the glass

will yield 18 parts of H20. If on the other hand the OH groups are associated with Na,

then on degassing the glass one mole of H20 will be produced for every one mole of OH,

viz., Na(OH) + H+ = H20 + Na+··· .. ·· '(20)

48

where a Na+ is exchanged for a H+ with the destruction of the Na(OH) ion pair. The above

mechanism depicted in equation (20) has recently been proposed by Kohn et. al., (1989) to

account for NMR data on hydrous albite glasses. The important point here is that pyrolysis

methods used to extract water from glasses eventually produce molecular H20 which is

then measured by some analytical techniques (e.g., gravimetric, mass spectrometric,

chromatographic) whereas infrared methods measure the concentration of X-OH groups

(where X = cation). Depending upon the relative importance of the above two mechanisms

in melts of different composition it would be reasonable to expect to find variations in the

£3550 em-I values if calibration techniques involve pyrolysis for the determination of total

water abundances. To test this hypotheses a combined infrared/mass spectrometric

technique was used to determine the value of £3550 em-I in two additional glasses with

low water contents and high silica (Table 4). The obsidian and tektite samples were chosen

since neither show bands in the lR due to molecular water and are silica rich compared to

the submarine volcanic glasses. The high silica contents should favor mechanism depicted

in equation (18) and yield extinction coefficients for the 3550 cm- I stretch larger than those

in Table 2 using the Beer-lambert Law and the pyrolysis/mass spectrometric water values.

This is indeed what we observe (e =200-205 L/mol-cm) and may help to rationalize the

variation in the e values reponed in Table 2. Further studies are, however, necessary using

a wider range of glass compositions to determine how critical major element chemistry and

total water are in affecting extinction coefficients.

The bending mode of liquid water appears at approxiamtely 1640 cm- I and is

estimated to have an extinction coefficient of 16 L/mol-cm (Thompson, 1965). The most

definitive evidence for the existence of molecular water in glasses would therefore be the

appearance of a band at - 1640 cm' '. The presence of a band at 1630 cm- 1 in the infrared

spectrum of an autoclaved thin film of Na2D-Al203-B203-Si02 led Ernsberger (1977) to

suggest that all the absorption at 1630 cm! can be accounted for if the adsorbed water is

49entirely dissolved as molecular H20 in the glass. Ernsberger's suggestion was based on the

assumption that the extinction coefficient of water in glass was the same as in liquid water.

Bartholomew et. al., (1980) have, however, calculated that the molecular absorptivity of

the 1635 cm- 1 bending mode of water molecules in a Na20-K20-Zno-Al2Si02 glass with

11 wt, % water as 28 ± 5 l/mol-cm, Silver and Stolper (1989) estimated the absorptivity of

the bending mode of water molecules in high-pressure quenched NaAISi30g-H20 glasses

to be 49 ± 2 l/mol-cm. Evidently the absorptivity of the bending mode of water molecules

depends on composition and perhaps on the mode of preparation of these glasses.

B. DISCUSSION

1. Effect of bulk composition on the speciation of water

The most interesting result of this study is shown in Figure 14a where the

absorbance data from Tables 2 and 3 were used to plot the ratio [Abs (1635 cm-1)/(Abs.

3550 em -1)] against total water content. Over the compositional range of glasses in the

present study this ratio is proportional to the ratio of molecular-to-total water. (If this were

not the case one would not obtain the linearity shown in Figure 13.). Only those glasses

for which we could confidently measure the absorbance (i.e., peak height) at 1635 cm- 1

for molecular water were used. For glasses with water contents less than approximately 0.5

wt. %, it was either not possible to reliably measure this band or it was not observed. It is

apparent that for these glasses at any given total water content there is no single value for

this ratio but rather a range of values. For example, for glasses with total water contents in

the range 0.5-1.0 wt, % this ratio may vary approximately lO-fold from 0.03 to 0.30. This

suggests that the relative proportions of the two hydrous species (i.e., molecular water and

X-OH groups) in these glasses is being inflenced by silicate chemistry. That this is likely

to be the case is demonstrated in Figures 14b and 14c which clearly show that increasing

the silica content and/or decreasing the concentration of non-tetrahedral cations (e.g., Na+)

causes the relative abundance of molecular water to increase. If one extrapolates the silica

content to 100 % in Figure 14b the ratio, Abs(l635)/Abs(3550), becomes almost 1. This is

expected since water can only exist in the molecular form in fully polmerized silica. A plot

50

0.4 r-----------------------.

oo

0.3 -

- 5tor<)tO

0tel!).::. !:2. 0.2 I- 0

00en en 0.0 .0« «

00

0.1 I-

0

0o 0Q)8s00

0

00I I

1.0 2.0 3.0H20 (wt.0/0)

Figure 14A. Relationship between the band intensity ratio for the molecular-to-total water

[Abs( 1635cm- 1)/Abs(3550cm- l)] and Total H20 (wt. %).

51

OAr---------------...---

0.3

--LOOr0LO<.OLO::: !2 0.2en en.c .c««

0.1

50 60Si 02 (wt, 0/0)

70

Figure 148. Relationship between the band intensity ratio for the molecular-to-total water

IAbs(1635cm- I)/Abs(3550cm- l )J and Si02 (wt.%).

52

0.4,..-------------------

o

0.4

oo

o

o 0

0.2 0.3Na20+K20+CaO+MgO ( 01<)

Si02+AI 203+P20S wt. 0

O'--------L..--------I-----~0.1

0.1

0.3

--L()OJ"{)L()<.OL().::::: !2 0.2en en.c .c««

Figure 14C. Relationship between the band intensity ratio for the molecular-to-total water

lAbst 1635cm- 1)/Abs(3550cm- 1)] and [(Na2D. + K20 + CaD + MgO)/(Si02 + A1203 +P2DS) I.

53similar to Figure 14c is obtained if one uses NBOrr (where NBO is the number of non-

bridging oxygens and T the number of tetrahedral cations; Mysen, 1988) in place of

[(Na20 + K20 + CaO + MgO)/(Si02 +A12~ + P205)] (Figure 14d). The correlation

coefficient is less, however, (0.717 vs. 0.821) which probably reflects the multiple

coordination of Fe (4 and 6-fold) in these glasses. Athough, there is considerable scatter in

the data points plotted in Figures 14b and 14c there can be little doubt that charge balancing

cations do play an important role in the dissolution mechanism(s) for water in these

petrologically important glasses.

2. Effect of quenching rate and annealing

The critical question from the point of view of constraining solution models of

water-bearing silicate melts is to what extent glasses preserve the concentration of hydroxyl

groups and molecular water of the melts from which they are formed by quenching. This

question will be very difficult to answer unequivocally without in situ spectroscopic

measurements of hydrous melts at high temperatures and pressures. There exists some

evidence (e.g., Aines et. aI., 1983) to suggest that hydrous glasses and melts are similar,

but much more work is needed in this area.

Kohn et. al., (1989) state that it is probable that some structural features of melts

will be more readily quenched into the corresponding glass structure than others (i.e., the

glass transition temperature for certain types of exchange will be different from the bulk

glass transition temperature). In particular, structural rearrangement between H20 and OH

in the melt could take place very rapidly compared with even the fastest achievable quench

rates.

It is possible that the variation in quenching rate of the magma on the seafloor is

responsible for the scatter observed in Figure 13. For example, my depth profile study on

the Loihi pillow basalt (see Appendix C), show that there is an apparent increase in the

proportion of molecular water from the more slowly cooled regions of the lava. It is

conceivable that the cooling rate of a silicate melt and hence the quench-induced stress on

the glass may playa role in determining the molecular water/OH ratios which are observed.

Previous calculations (Bonatti and Harrison, 1989) have shown that for submarine basaltic

lavas of the same composition, the intensity of the stresses induced in the cooling lavas

54

OAr-------------------.

oo

0.3

LOar0 lDtOlD- ro 0.2--(/) (/)

-0 -0<{<{

0.1

0.6NBO

T

0.8

o

Figure 14D. Relationship between the band intensity ratio for the mokcular-to-total water

Il\b~(I635ct11-1 l/Ahs(35S0cm- 1)I and NBOrr.

55decreases with rate of discharge of the lava at eruption. This results in pillow flows being

under higher stress than lavas producing sheet flows. Dingwell and Webb (l989a,b) have

suggested that for rhyolitic melts water speciation should be temperature dependent, and,

Stolper (1989) has shown that for rhyolitic glasses the ratio of water molecules to hydroxyl

groups can be altered by changing the synthesis temperature. Although the effect is not

large this ratio is increased by lowering the temperature of heat treatment. However, since

most submarine magmas are believed to erupt at similar temperatures (from between 1100

to 12000C for basalts, Rodey Batiza. private communication) and onto the seafloor where

the mean temperature is -40 C worldwide, it could be argued, that at least for glasses

sampled from the outermost rind of quenched lavas, the effects of slight differences in

temperature and cooling histories on water speciation would be small.

Mysen (1990) has also noted that the spectra obtained from temperature-quenched

melts (hydrous glass) could be misinterpreted since the proportion of molecular/hydroxyl

water may be dependent upon the conditions of temperature quenching.

Stolper and coworkers in several reports (e.g., Stolper, 1982a; Silver and Stolper,

1989) suggested that the water speciation is only weakly temperature dependent. On the

other hand Dingwell and Webb (1989a,b) have suggested that speciation in silicate melts in

general, and OH-/H20 in particular, was significantly temperature-dependent. Dingwell

and Webb (l989b) calculated from relaxation theory, for example, that for a reaction:

20H- = H20 + 0 2- · .. ·(21)

which could describe the equilibrium between oxygen in the melt (02-) and the two

principal forms of water, ~H = 25 ± 5 kJ mol"1. They concluded that the OH-/H20 ratio

measured from quenched materials is that which existed in the glass when frozen in near

the glass transition temperature, and that at the temperature of equilibration, nearly all the

dissolved water exists in the OH- form. In contrast, Silver and Stolper (1989) suggested

that reaction (25) was only weakly temperature-dependent.

56The use of multinuclear magnetic resonance can be beneficial here because it yields

a much more complete picture of the structure of the glass than any other single technique

by providing information about the spatial relationship between atoms of different types

(except x-ray crystallography). It can help provide clues as to which features are most

likely to have arisen during the quench. This is in contrast to previous infrared

measurements which have concentrated solely on water speciation.

3. Comparison with previous studies

Most previous investigations on the nature of water in silicate glasses have

examined specially prepared synthetic samples of relatively simple composition and high

water contents (Newman et. al., 1986 and Silver and Stolper, 1989). For example, Silver

and Stolper (1989) have used infrared spectroscopy to measure the abundances of

molecular water and hydroxyl species in high pressure/temperature quench melts of albite

compositions with total water contents in the approximate range 1-10 wt, %. In contrast to

these, the samples used in the present study were all naturally-occurring glasses of basaltic­

dacite composition with water contents 0.10 to 1.738 wt.%. Due to the more complex

nature of our glasses a strict comparison of previous observations with our results is not

possible. However, Silver et. al., (1990) showed for some rhyolitic glasses that those

which were synthesized at lower quenching rates there was an increase in the ratio of

molecular water to OH groups compared to those which were quenched more rapidly

suggesting that OH groups convert to molecular water as the melt cooled. This is consistent

with our observation that glass taken from the more slowly cooled deeper layers of the

glassy rind of pillow basalt showed higher ratios of molecular-to-total water than that taken

from the more rapidly quenched outermost layers (see Appendix C).

Previously, hydrogen manometry or 1H NMR spectroscopy were used to obtain

total dissolved water contents (Newman et. al., 1986 and Silver, 1988). Here, we have

used high temperature mass spectrometry. Newman et. aI., (1986) emphasized that

only the extinction coefficients for the 4500 ern-1 and 5200 cm- 1 bands were determined

directly from the manometric measurements of total water content. This study determined

57

the extinction coefficient for the OH-stretch band at 3550 em -1 directly by using the total

water determined by HTMS. In this study however the near-ir bands observed by

Newman et. al., (1986), were not observed due to lower total water contents.

The focus of most previous work has been directed to the relationship between total

water content and the concentration of molecular water and hydroxyl groups (Bartholomew

et. aI., 1980; Acocella et. al., 1984; Silver and Stolper, 1989). The basic conclusion from

these studies is that at low water contents hydroxyl groups are the dominant hydrous

species but as total dissolved water exceeds a few weight percent additional water dissolves

as molecular water. The results of the present study are in qualitative agreement with this

behavior since we could not either detect or reliably measure the presence of molecular

water in most glasses with total water contents less than 0.5 wt, %.

The influence of bulk composition on the speciation of water in silicate glasses has,

however, been addressed in one previous study. Silver et. aI., (1990) have used infrared

spectroscopy to measure the abundances of molecular water and hydroxyl species in

synthetic rhyolite, orthoclase, jadeite and Ca-AI silicate glasses with water contents from 1

to 12 wt.%. Athough the trends in the species concentrations in all their compositions were

similar (as described above) it was observed at similar total water contents that increasing

silica and K relative to Na favor molecular water over hydroxyl groups.

For basaltic glasses, similar to those studied here, Dixon et. aI., (1988) obtained a

value of 63 ± 5 Lrnol-cm which is within the experimental uncertainty of our value of 61 +

1 Lmcl-cm. Although we used different methods for total water determination this appears

to indicate that the two methods are comparable in their ability to determine total water in a

glass sample. It is interesting to note however that the stretching band envelope for liquid

water has a value of 81 L/mol-cm (Thompson, 1965).

Recent NMR studies on albitic glasses Kohn et. aI., (1989) has also placed the

dissolution mechanism of water in glasses in question. Spectra for 29Si and 27Al showed

little change throughout the range of water concentrations studied (0-67 moI.%).

However, major changes were observed in the 23Na environment. They concluded that

58water produced no depolymerization of the silicate network and that hydroxyl is stored as

NaOH complexes. These data are consistent with previous Raman data (Mysen and Virgo,

1986a) which indicated the presence of molecular water and no Si-OH or AI(4)-OH

groups. The hydroxyl group was suggested to be associated with network modifying six-

fold coordinate A13+ or Na + ions. Okuno et. aI., (1987) using x-ray radial distribution

analyses on synthetic albitic glasses, also observed no drastic change in the TO 4 tetrahedra.

If glasses with more complex chemistry like those in the present study are shown to behave

in a similar manner existing models for the effect of water on silicate melt structure must be

re-eval uated.

C. Future Work

This study used naturally occurring silicate glass samples with a range of

compositions and perhaps different quenching histories. It is neccessary to do an in-depth

study where the composition is fixed while the volatile content is varied. This wasn't

possible here since too many variables had to be considered. Future work should be

directed to the preparation of synthetic high temperature/pressure glasses for a better control

upon composition and quenching rates. The quenching rate could be systematically varied

to see how this effects the speciation of water in the glass. These glasses could then be

degassed with the high temperature mass spectrometer and studied with infrared

spectroscopy. The near-ir bands would be used for compositions where the iron ion content

was low or zero. Also, to check how silica content affect the extinction coefficient a

systematic study could be undertaken where only the silica content varied for a fixed

amount of total water.

59APPENDIX A

Obsidian (An interlaboratorycomparison)

In order to evaluate the variability in FfIR data between different laboratories,

using different spectrometers, an interlaboratory comparison study was undertaken with

Dr. Sally Newman (California Institute of Technology).

This study involved obtaining the infrared spectra in the near-ir region of five

obsidian glasses (ave. Si02, 76.5 wt.%) from the ca. 1340 A.D. eruption of the Mono

Craters chain in central California (Newman et. al., 1988). Five doubly polished glasses

were provided. From previous infrared and monometric studies the extinction coefficients

for the 4500 and 5200 cm- l bands were known. These are the combination bands for

hydroxyl and molecular water respectively. The densities for all the glasses were assumed

to be 2277 gIL, a reasonable value based upon comparison with other glasses of similar

major element chemistry. Thickness were measured by a thickness monitor <± 2 urn).

Figure 15 shows the near-ir spectra of two glasses with different amounts of total

water content. These spectra were used to calculate the amount of hydroxyl and molecular

water present in the glasses. The sum provided the total amount of water in each glass.

This is only true if one assumes that there are only two ways in which water can exist

within silicate melts. This was shown to be a good assumption because in their study

when they plotted the total water obtained by infrared spectroscopy against the total water

obtained by monometric analyses a linear relationship was obtained with a good correlation

coefficient (Newman et. aI., 1986). Table 5 compares my results with values obtained by

Caltech. As can be seen there is good agreement between the two laboratories except for

sample 8a. This demonstrates that two different spectrometers can provide reproducible

water values. This represented the first opportunity to check how analyses from this

laboratory compared with other groups. The final results of this study are still not available.

In addition to high total water contents, these glasses also had dissolved molecular

carbon dioxide within them. This is indicated by the presence of a band at -2350 em -1 due

to the antisymmetric stretch of dissolved C02' Figure 16 shows the spectra of two glasses

~C)

~Cd

..05-0oen~

<

.... 52()() cxr I (MOLECULAR H20)"COMBINATION (\)1 + \)2)

J.

i~IC84-bb-4a # 5a (2.66 wt. % H20)

Thickness =635 J.l111

t-vICg4-bb-4a # 9a (2.72 W£.% H:{))

Thickness = 1024 urn

\

r'~5(lO C i\I- I (0 II )

COMBINATION (lJl + tJ(T-O.T»

5000 4000Wavenurnbers (cm e- t )

Figure 15. Infrared spectra of two water rich obsidian glasses from the ca. 1340 AD Mono

Craters eruption showing the near-infrared bands at -5200 cm- 1(H20) and -4500 cl11- 1

(X-OH) in the 6000-4000 cm- 1 region.~

Table 5. Infrared data in the near-infrared re~ion for ~Iasses from theeruption of ca. 1340 AD of the Mono craters chain in central California

Sample # Thickness 4500 cm- 1 5200 cm- 1

MC84-bb-4a (1lt11) (wt.% OH) (wt.% H2O)

4a 336 0.42(0.41) 0.07 (0'{>7)

Sa 635 1.19( 1.26) 1.47( 1A5)

6a 800 0.41 (0.40) 0.07(0.06)

Sa 118 1.06(0.96) 0.92(0.74)

9a 1024 1.24(1.26) 1.48( 1.45)

Bracketed values are those obtained at Caltech.

61

(1)Co)

=C1S.0J-.oIl2:a

II MOLECULAR C02 IN OBSIDIAN GLASSES!

1lII

.5

II 2300~ 24'00 bers (cm.-1)WavenuID.

Figure 16. Infrared spectra of two water rich obsidian glasses from the CA. 1340 ADMono Craters eruption showing the presence of dissolved C02 in the glasses in the 2500-

2200cm-1 region.

i2200

0­N

63that contained dissolved molecular C02' Using their extinction coefficient (of 945 L/mol-

em) it was possible to calculate the amounts of dissolved C02 in the two glasses. Since the

glasses studied here were not specifically mentioned in their paper (Newman et. al., 1988)

it was not possible to check the accuracy of the data. However, values are close to some

of the glasses quoted in the paper. Since these glasses probably contain minor amounts of

carbonate (C032-) the total dissolved carbon dioxide in the glasses is unknown.

64APPENDIX B

The speciation of water within hydrous alkali silicate glasses

Much work has been carried out to study the speciation of water within high­

temperature/pressure glasses. For example, hydrous albitic glasses have been synthesized

at pressures between 8 and 25 kbar and temperatures between 1000 and 16000C (Silver,

1988). It has been shown that at high temperatures water is primarily present as hydroxyl

groups for water concentrations up to 4 wt, %. There is, however, very little work to

understand the speciation of water resulting from low temperature alteration of volcanic

glasses.

The ages of subaerial volcanic glasses are often determined by K-Ar dating (Hall et.

aI., 1984). Volcanic glass is readily altered to more stable phases, generally smectites or

other clay minerals. Accompanying this mineralogical change are many chemical changes,

such that the original chemical composition of the glass cannot be discerned in the final

product. Cerling et. a!., (1985) studied the alteration of subaerial glasses and noted that

significant amounts of K and Na are lost from the siliceous glasses. Since Ar is much less

mobile relative to K, will result in discrepant K-Ar ages for glass with sigificant hydration.

This loss has been attributed to an ion exchange process where there occurs significant

hydrogen exchange for alkalis (Ernsberger, 1980). Cerling et. aI., (1985) noted the low

temperature hydration results in molecular water being dominant over hydroxyl groups.

This observation differs from that of the high-temperature water speciation where hydroxyl

groups are observed to dominate in the glass.

In-situ Raman studies by Exarhos and Conaway (1983) on the dissolution of glass

in aqueous media suggested that leaching occurs via ion exchange reactions between

sodium cations localized in ionic glass sites and hydrogen ions from the solution.

However, Ernsberger (1980) proposed that free protons cannot exist in the glass and must

be transported as hydronium ions (H30+). Tsong et. al., (1981) used nuclear reaction

techniques to show that the number of hydrogen atoms in a leached glass is three times that

65

of sodium atoms, implying that the interdiffusing species are indeed hydronium (H30+)

and sodium ions.

In order to understand the hydration of volcanic glasses at ambient (or low

temperature), a combination of infrared, Raman and magic angle spinning nuclear magnetic

resonance (MAS NMR) spectroscopy were used to study the speciation of water within

synthetic alkali silicate glasses with the compositions: 30%Li20-70%Si02' 45% Na2Q­

55% Si02 and 35 % K20-65 % Si02' These glasses which were prepared by the National

Bureau of Standards (NBS), had hydrated over a period of -6 years at ambient conditions

in the spectroscopic laboartory of the Hawaii Institute of Geophysics (RIG). This provided

a unique opportunity to study the speciation of water in some simple alkali silicate glasses

which had hydrated at ambient conditions opposed to being hydrated at high temperatures

and pressures, as is usually the case. The lithium glass appeared the least hydrated when

handled, whereas, the sodium and potassium glasses, were close to gels.

Near-IR spectra for the K and Na glasses displayed strong bands at - 5200 cm- 1,

which are attributed to the presence of molecular water in the glasses (Newman et. al.,

1986; Figure 17a,b). However, no bands were observed at -4500 em-I, which means that

no hydroxyl groups (other than bonded to hydrogen in molecular H20) were present or

that the concentration was so low that we couldn't detect them. Since the thickness ofNa

glass (-900 urn) was much greater than that for K glass (-200 urn) it was obvious that the

K glass has more water. The lithium glass showed no bands in the near-ir, although a very

weak band at - 3550 crn! was observed (Figure 18). This means that the concentration of

total water is extremely small, especially since the sample thickness was -715 urn.

Raman spectra for the Na and K glasses were obtained using the micro-Raman

facility in HIG (see Experimental Methods section, Part 2). The Raman spectra for the

K20-Si02 and Na20-Si02 glasses have strong bands in the 3000-4000 cm- 1 region. Their

spectra which were very similar to that for pure water can be deconvoluted into three

distinct bands (Figure 19). Brooker et. al., (1989) from their Raman study on water

66

HYDRATED 35%KZO-65%SiOZSYNTHETIC GLASSd-ZOOllm

-5200 CM-l (MOLECULAR H20)

HYDRATED 45%NaZO-55%SiOZSYNTHETIC GLASSd-900J.U11

o~ . _

~ 2UZ<=~oen 1=<

-5200 CM-l (MOLECULAR "20)

o ..L.- ._--~------~

5000 4000

WAVENUMBERS (eM-I)figure 17. Infrared spectra for NaZO and KZO-SiOZ synthetic glasses in the near-ir region

(6000-4()()() cm- 1) .

67

GLASS OF THICKNESS ....715J.lrn

.4

-3550 CM-l

U(OH) STRETCH (H20 + OH)

3000 2000

WAVENUMBERS(CM-1)

o .l.--------.,.....-----------:-:i

Figure 18. Infrared spectrum of a hydrated Li20 -Si0 2 synthetic glass in the region 4000­

2000 em-I.

68

3800

HYDRATEDK20-Si02

GLASS

3200 3400 3600

RAMAN SH1FT (em")3000

i>-I--

HYDRATED-(f)

Z No2O-Si0 2wI--Z

<.9Z-0:::WI--I--«u(f)

H20 20°C

Figure 19. Deconvoluted Raman spectra for water and hydrated Na20- and K20-Si02

glasses in the 30()O-3~()() em I region.

69

containing 160 and 180 isotopes, attributed these bands to the following three types of

water interactions: symmetric complex or polymeric water (-3252 em-1), asymmetric or

dimeric water (-3384 cm-1) and almost interaction free-water (-3620 cm-1). It is noted

from the relative normalized areas that the amount of interaction-free water increased from

-3% for pure water to -9% for the two glasses. Whereas, the amount of asymmetric water

decreased from -67% (pure water) to -62% and -53% for the Na- and K-silicate glasses,

respectively. The amount of polymeric water increased from -30% for pure water to -38%

for Na-silicate glass whereas, the value decreased slightly to -29% for K-silicate glass.

A band at - 1652 cm-1, corresponding to the bending mode of molecular water,

was observed for the K glass (Figure 20). Relative to pure water (-1637 em -1) this band

appears at higher frequency. It has been noted that the bending mode of water is at -1594

em-I for water in the vapor phase whereas this band appears at -1645 cm- 1 for water in

the liquid phase (Fernandez-Prini et. al., 1992). The shift to higher frequency observed in

the bending mode band of water when vapor condenses is due to hydrogen bond

formation since the existence of an O---H-O bridge will stiffen the bending mode because

hydrogen bonding favors linearity in the O---H-O hydrogen bonded arrangement of atoms,

thus blue shifting the band. Raman bands that could be attributed to the presence of Si-OH

bonds (in 900 to 1000 cm- 1 region) were not observed in the spectra of the K and Na

hydrated glasses. Since the lithium glass has a very low water content, as indicated by

infrared spectroscopy (due to its higher sensitivity), no bands that would indicate the

presence of water were observed in the Raman spectrum.

A MAS nmr experiment was carried out on these glasses by Dr. Barbara Sherriff of

the University of Manitoba, Winnipeg, Canada at the Prairie National Research Centre.

Both the hydrated glasses and samples of the same glasses dehydrated by remelting were

sent to her to examine the effect of water on their structures.

The 29Si spectra (Figure 21 and 22) for the Na and K glasses gave two peaks: Na

glass hydrated at -76, -87 ppm; remelted at -77, -84 ppm and K glass hydrated at -86,-95

70

1---(A) BENDING MODE OF WATER

FOR PURE WATER

1

(B) BENDING MODE OF WATER TFOR 35%K20-65%Si02 GLASS

i

1600 1800

RAMAN SHIFT (eM-!)

Figure 2(). Raman spectra of water and hydrated K20-Si02synthetic glass in the 1500­

2000 cm- 1 region.....

00wr----:1'-00I II I

(A) HYDRATEDNa20-Si02 GLASS

00r--:~1'-00I I

I I(B) DEHYDRATED

Na20-Si02 GLASS

71

-20 -60 -100

PPM

-140 -180

Figure 21. 29Si MAS NMR spectra for hydrated and dehydrated Na20-Si02 glasses.

72

-20 -60 -100 -140 -160

PPl\l

Figure 22. 29Si MAS NMR spectrum for hydrated K20-Si02 glass.

73

ppm. The 29Si NMR spectrum for the remelted K glass wasn't possible because it was

difficult to pack the rotar evenly for this glass making it difficult to get it to spin. A static

spectrum (without spinning) was tried for this glass but gave poor results. The Li

glasses gave a single similar symmetrical broad peak: hydrated at -93 ppm and remelted at

-91 ppm (Figure 23a,b). From previous work on gels (Personal Commun. B. L. Sherriff),

it was found that there is a shift of about 8 ppm to high field with each additional Si-O-Si

bond. The following is a tentative assignment: -72 ppm (00), -79 ppm (Q 1)' -87 ppm

(Q2)' and -95 ppm (Q3)' The Q number refers to the number of Si-O-Si bonds. These

glasses exhibit differences due to changes in the degree of polymerization with an order of

increasing polmerization of Na < K < Li, as might be expected by their respective

compositions (45% Na20-55% Si02,35 % K20-65 % Si02 and 30%Li20-70%Si02)'

The slight increase in intensity of the low field peak (Q 1) in the 29Si spectra of the

hydrated Na silicate glass can be interpreted as a very small amount of depolymerization

(such information is not possible from IR measurements). The only difference between the

spectra of the hydrated and dehydrated Li glasses was a small shift to high field in the 29Si

spectra (Figure 23a,b) which could be due to shielding by water molecules associating with

the Si04 tetrahedra. There are also smaller changes in chemical shift caused by

electronegativity differences of the alkali cations. However, these are usually about 1 ppm,

which is of the same order of magnitude as the error in the measurement of chemical shifts

of these broad peaks from the glasses.

The difference between the 23Na spectra for the hydrated (-3 ppm) and the

dehydrated (-5 ppm) glasses (Figure 24a,b) can be attributed to Na coordinating with H20

molecules in the hydrated glass. This peak appears at higher field due to the greater

shielding of the Na+ cation. The sharp peak at -7 ppm in the remelted Na glass could be

due to Na+ as it is similar to Na+ in NaCI (Sherriff et. aI., 1987). The difference in shape

of the major peak is probably due to different quadrupolar parameters of the 23Na nucleus,

which is related to the symmetry of the site.

oroen

I

I

enI

I

(A) HYDRATEDLi20 -Si0 2 GLASS

74

(B) DEHYDRATEDLi20-Si02 GLASS

-20 -60 -100

PPM

-140 -180

Figure 23. 29Si MAS NMR spectra for hydrated and dehydrated Li20-Si02 glasses.

'<---~

or()I

I

oq~lO

I

I

(A) HYDRATEDNa20-Si02 GLASS

(B) DEHYDRATEDNa20-Si02 GLASS

75

80 40I

o

PPM

I

-40i

-80

Figure 24. 23Na MAS NMR spectra for hydrated and dehyclratedNa20-Si02 glasses.

76

oI

I (A) HYDRATEDLiZO-SiOZGLASS

oI

I(B) DEHYDRATED

LiZO-SiOZGLASS

I

10I

5I

oPPl\1

i

-5i

-10

Figure 25. 7Li MAS NMR spectra for hydrated and dehydrated Li20-Si02 glasses.

77

The 7Li spectra for the hydrated and dehydrated samples were identical, with one

broad peak at --0.1 ppm (Figure 25a,b). This is in aggrement with the infrared spectrum

which indicated that the hydrated glass contained very little water.

From these results it is possible to deduce that the hydration of these synthetic alkali

silicate glasses increase as the size of the cation increases. These observations are

consistent with the observations of Matson et. al., (1983) who noted that the Raman

frequency of the Si-O- (non-bridging) vibration was lowest for the lithium-silicate glass

and increased as the size of the cation increased. This means that the non-bridging oxygen

is strongly bound to the lithium cation and that the bonding strength decreases as the size of

the cation increases. It is, therefore, more difficult to hydrate the lithium glass but the

hydration gets easier as the size of the cation increases. Unlike high temperature/pressure

glasses the water is principly present as molecular water as clearly verified by the infrared

and Raman spectral data since we observed no Si-O-H stretching bands. Very little

depolymerization of the hydrated alkali silicate glass has taken place as indicated by the

MAS NMR data. The difference in depolymerization observed can be attributed to the

percentage alkali oxide present in the glass. The Raman spectra in the 3000-4000 cm-1

region of the hydrated Na- and K-silicate glasses are similar to that for pure water and

suggest that the water molecules are engaged in three different types of interactions. The

relative amounts of each type appears to be effected by the nature of cation in the glass.

In conclusion, infrared, Raman and MAS NMR spectroscopies are useful tools for

studying the ambient temperature and pressure speciation of water in glasses. Further

studies of hydrated glasses of a variety of compositions are expected to enhance our

understanding of the alteration of volcanic glasses.

78APPENDIXD

The effect QfQuenching rate and temperature Qn the mQlecular-to-tQtal water ratiQ

1. Depth prQfile study Qf the glassy rind from Loihi seamQunt

The glassy rind formed on submarine pillow basalts when lavas are erupted at

depths greater than approximately 1200 m is known tQ retain most of the original volatiles

present in the magma (MQQre, 1970). The outer glassy rind results from the rapid

quenching experienced by the extruded lava (lIOO-1200Qe ) at the bottom of the seafloor

(-4QC). Except for a few microcracks this layer effectively isolates the slowly-cooling

interior from the Qcean environment. An important question arising from this process is,

how does the speciation of water change as a function of quenching history? This question

hasn't been addressed previously, Therefore, a combination of mass spectrometry and

FTIR spectroscopy were used to study this problem.

TQ investigate this problem a sample from the summit of Loihi seamount, Hawaii's

youngest active volcano, located 50 km south of Kilauea's summit and approximately

lOOOm below sea level, was chosen. This sample was selected because it had a thick rind

of fresh, alteration free glass and had previously been studied for volatiles and major

elements (Garcia et. aI., 1989). Table 6A and 6B summarizes these data. A schematic

cross-section of a typical pillow basalt when viewed with a binocular microscope, can be

divided into four distinct regions (Figure 26) as described previously (Aggrey, 1989).

(l) The outermost region, here identified as region 1, was about 4 mm thick and consisted

of pale-brown transparent glass virtually free of phenocrysts.

(2) Region 2 (also approximately 4 mm thick) was mostly glass, similar to that in region 1,

but appeared darker and contained grey specks.

(3) Region 3 stretched Qver a depth of about 10 mm below the lower boundary of region 2

and consisted of a mixture of small amounts of glassy material, grey specks (as seen in

region 2) and non-vitreous rock.

(4) Region 4 (beIQW region 3) was grey rock with virtually no vitreous material.

Three regions were clearly visible in thin-section with regions 3 and 4 merging together in a

79

1hhlc 6A. Loihi Deplh Profile Sludy: Volalile Abundances (wt.%J

SAMPLE "2° CO2 S CI F

I..J\YER I O.5RX 0.033 0.152 (),()95 0.006

LAYER 2 0.546 0.035 0.144- 0.099 0.004

lAYER 3 O.74X 0.026 0.151 0.112 0.004

I.AYER 4 0.953 0.039 0.141 0.109 0.003(Irom Aggrcy, IlJX9)

Tithle 613. Microprobe analvsis or Loihi J -4 glass (Laver 1) (Wl.r/r,)

4X.(i9 2.R2 13.4:1 11.X2 (i.X1 11.6R 2.54 0.54 0.32 9R.(i6

(Iron: Garcia C1. al., 19X9)

80

Transparent brown glassA1

Speckled brown glassA2

Black opaque "glass"A3

Aphanitic basaltR4

o, 1!

2 CENTIMETERS,

Figure 26. Composite sketch of the outer portion of a submarine pillow fragment. R1, R2.R3 and R4 refer to regions 1,2, 3 and 4 as discussed in the text (modified after Moore,1966).

81diffuse boundary. There was a slight decrease in the measured water content in going from

layer 1 (0.588 wt.%) to layer 2 (0.546 wt.%). However, these values can be considered

to be essentially identical since they are within experimental error. The water content for

layer's 3 (0.748 wt%) and 4 (0.953 wt%) increased markedly. In addition to differences in

water abundances significant differences were also observed in the H20 release profiles

(Figures 27 and 28; Aggrey, 1989). Layer 1 showed the typical bimodal release pattern

observed for alteration free submarine basaltic glasses taken from the outer most rind.

However, the release profile for layer 2 showed a shift in the position of the high

temperature peak to lower temperature with possibility of the low temperature peak being

masked by the high temperature peak. The release behaviour for layer 3, however, showed

two significant differences in comparison to layers 1 and 2. These were: (1) a shift of the

high temperature H20-release peak to lower temperature and, (2) the low temperature

"peak" was either buried or replaced by a large, broad envelope of H20 release beginning

at approximately 250 0C. This is consistent with the large increase in total water content

(Table 6A ) over those in layers 1 and 2. Layer 4 showed the same broad low-temperature

H20 release band (250-750°C) observed from layer 3 (Figure 28). The low-temperature

peak observed in layer 1 was no longer visible, but maybe it is buried under the sharp high­

intensity peak centered around 7500C. This sharp peak does not appear to be due to the

same process giving rise to the high-temperature peak from the glassy samples (layer 1 and

2). The intensity and sharpness of this peak is most likely due to a non-glassy portion of

the sample, suggesting the decomposition of a hydrous phase.

To study these samples with FTIR spectroscopy doubly polished thin-sections were

prepared and thicknesses measured as explained previously in the experimental section (see

Experimental Methods section Part 2). The Perkin Elmer 1720X FfIR was used for all IR

measurements. In Table 7, the thickness-corrected ratio of the peak intensity of the

molecular water band I(1635) to the intensity of the total water band 1(3550) is tabulated.

From Figure 29, which shows the infrared spectra for layers 1-3, it can be seen that the

amount of molecular to total water increases both as a function of total water content and

1.20

SAMPLE SIZE = 60.645MG HEATING RATE = 5.71 oC/MIN

LAYER 1

82

50e,

~ 0.80enzUJI-Z

Z0 0.40

0.0035. 435. 835. 1235.

TEMPERATURE 0cSAMPLE SIZE = 63.960MG HEATING RATE = 5.71 oC/M1N

1.20

LAYER 2

5o~

~en 0.80zUJ!Z.zo

0.40

0.00

35. 435. 835.

TEMPERATURE 0c1235.

LOIHI 1-4 DEPTH PROFILE

Figure 27. HTMS H20 release profiles for Loihi (LO-4) layers I and i. (after Aggrey,1989).

RATE = 5.72oC/MIN 83SAMPLE SIZE = 64.680MG HEATING

" -- H2O+ LAYER 3

1.20

~0e,

~ 0.80(f)

zUJI-Z

z

~./00.40

0.00

35. 435. 835. 1235.

TEMPERATURE °cSAMPLE SIZE = 64.020MG HEATING RATE = 5.71 oC/MIN

LAYER 4

1.20

~enzUJI­Z

Zo

0.80

0.40

0.0035. 435. 835.

TEMPERATURE -eLOIHI 1-4 DEPTH PROFILE

1235.

Figure 28. 'ITMS "10 release profiles for Loihi (1..0-4) layers 3 and 4. (after Aggrey,1989).

Table 7. The ratio I(1635)/I(355Q) for Loihi Depth profile study

LAYER #

2

3

4

1(1635)/1(3550)

0.0423

0.1462

0.2282

0.3244

00~

cuc:.>s::es

..0s...oen

..0<

LOIHI DEPTH PROFILE STUDY

4000 3000 2000Wavenumbers (ern-i)

Figure 29. Infrared spectra for Loihi (LO-4) layers 1, 2 and 3 in the 1200-4000 cm!region. 00

IJl

86depth below the glassy rind from which the sample was taken. The spectrum for layer 4

was very weak due to the poor quality of the glass. The slower cooled interior portions

have more molecular water. These results clearly show that the speciation of water is a

function of the thermal history of the sample. Hence, hydroxyl groups in the more

slowly cooled and partially crystalline interior apparently prefer to combine to generate

molecular water. A similar observation was made by Silver et. al., (1990) who studied the

effect of slow (3 °C/sec) versus fast (400 °C/sec) quenching rates upon the relative

amounts of molecular water and hydroxyl groups. They noted that for those glasses (of

rhyolite, orthoclasse and jadeite compositions) synthesized at slow quenching rates the

amount of molecular water was higher relative to fast quenched glasses.

2. The effect of annealing upon the ratio of molecular-to-total water within a volcanic glass

The glass sample selected for this study (MA-21) was taken from Mahukona, a

small submerged shield volcano located off the northwest coast of the island of Hawaii. Its

summit is approximately 1200 m below sea level (Garcia et, al., 1990). The composition of

this glass is weakly alkalic, similar to that from Loihi seamount. Chips of this glass sample

were annealed at - 580 °c for Ihr, 2 hr, 3 hr, 6 hr and 16 hrs. This temperature represents

the minimum temperature just before the first low temperature water release as displayed on

the mass pyrogram for the sample (Figure 30).

An isolated tube furnace (Daltech INC. Model DT-31-VT-05) was utilized for this

study. The furnace was pre-set at -580 °c before each sample was introduced. The sample

chips were held within alumina liners, normally used for mass spectrometric analyses. The

temperature was monitored using a Pt-Pt lO%Rh thermocouple. Before placing the sample

within the hottest part of the tube furnace, oxygen was purged from within the furnace by

using a constant flow of N2 gas for -10 mins. Then the sample was carefully lowered and

left there for the appropriate amount of time. Once the sample had been annealed, it was

carefully pulled out of the hottest region and allowed to cool in a cooler portion of the

furnace. It was carefully examined for any visible changes using a microscope. The only

~.""'"",.,..,~

.-.--.--.

1.00~

I I I

H2O+I

IJ'0.80

~u.J~;:)

~O.60~

0:::t.lJ,..-~L.:.J

~[-{

~

0.40

0.20

0' 1 " I ,I!!!! I ! I ! It! I

450 850 1250TEMP (OC)

Figure 30. H20 release profile from HTMS for Mahukona (MA-21) to show the annealingtemperature. ::x:

--.)

88change observed was a dull appearance on the surface. These chips were subsequently

doubly polished for infrared analyses as discussed in the Experimental Methods section

(Part C (2». The Perkin Elmer 1720X Ff-IR was utilized to study the sample in the 400-

4000 cm! region. The sample was examined at a few points with the aid of a 600 J.1M

aperture. The two peaks of interest are due to the bending mode of molecular water at

1635cm-l, and the fundamental X-(OH) stretch.where X can be Si or Al observed at 3550

em- l (Figure 31).

The ratio of the thickness-corrected peak intensities 1(1635)/1(3550) were calculated

for both the unheated and the annealed samples. As seen in Table 8 there appears to be an

increase in this ratio with increased annealing time. There is a marked change in this ratio

when the sample was annealed for only 2 hours. This ratio clearly shows an increase in the

amount of molecular to total water. This is not entirely unexpected since we observe the

release of water as molecular water (H20) during the mass spectrometric degassing

analyses. Consequently, the hydroxyl groups present in the glass must combine to form

H20 during the annealing process. This process is just the reverse of the reaction that

occurs when water is sited into the silicate melt as it is quenched on the seafloor. The drop

in the 1(1635)/1(3550) ratio at 3 hrs may be attributed to the loss in molecular water from

the glass. This ratio increases for 6 hrs as further hydroxyl groups convert to molecular

water.

(1)c:,)

~~

..cJ...oen.0<

ANNEALING STUDY ON MAHUKONA GLASS MA-21

1 hr.

2

3

o ,4000 3000 2000

Wavenurnbers (em-1)

Figure 31. Infrared spectra of Mahukona (MA-21) glass for unheated and samples.annealed for Ihr, 2 hr and 3 hr, respectively.

oc-c

Table 8. The relationship between the 1(1635)(1(3550) ratio with annealing time.

Annealing time I(1635)!I(3550)

unheated 0.0549

1 hr 0.0724

2 hr 0.0839

3 hr 0.0615

6 hr 0.0723

16 hr 0.0850

\Do

91APPENDIXD

Water-poorand water-rich glasses

1. Water-poor glasses

Four water-poor glasses were analyzed for water using high-temperature mass

spectrometry (HTMS), Ff-infrared spectroscopy and probed for major elements. This

study was undertaken to investigate what effect composition, quenching rate and low water

contents have upon the molar absorptivity coefficient. Two of the glasses are from

subaerially erupted lavas, one is a tektite and the other is an obsidian.

Elemental analyses were carried out using the fully automated 3-spectrometer RIG

Cameca microprobe operated in a wavelength dispersive mode. Probe analyses were made

on a carbon coated resin plug that contained individual grains representative of each sample

used for mass spectrometric and infrared analyses. The glasses used as standards (and the

elements for which they were used) were Makaopuhi ( Ti, Mg, Ca, and Fe) and VG-568

(Cameca standard) (Na, K, AI, and Si). A beam diameter of -5 urn was utilized. A 15 lev

filament voltage and a counting time of 10 seconds per element was used in all cases. Raw

data were adjusted on-line for detector dead time and element background, and analyses

were reduced using ZAF corrections. Each reported analyses represents the average of at

least 3-4 spots per grain and two or more grains per sample.

The high temperature mass spectrometric analyses were carried out as outlined

previously in the experimental section. However, due to the low volatile content of these

glasses larger sample sizes (-75 mg) were necessary to allow for a more reliable wt.%

H20 measurement.

The densities were measured with a pycnometer. Doubly polished thin-sections of

each of the samples were prepared using the technique outlined in the experimental section.

The platlets were made thicker than normal since the water content was much smaller than

those for submarine basaltic glasses.

Elemental analyses, density, thickness, infrared absorbance and extinction

coefficient (see Experimental Methods section for calculations) for each sample are given in

Table 9.

Table 9. Major elemental and infrared data for low water content ~lasses

Sample Obsidian Tektite Kalapana * Kilauea Iki **Si02 76.93 72.83

Ti02 0.05 0.7

Al203 12.51 11.91

FeO 0.84 4.52

["lgO 0.02 1.8

CaO 0.47 1.64

Na20 3.89 1.25

K20 4.59 2.56

H20 0.073 0.006

Density (gIL) 2436 2421

Thickness (urn) 267, 159 506

Absorbance 0.518, 0.321 0.081

Abs. Coeff. # 100. 205 201(l/rnol-cm)

50.51

2.39

13.6

10.65

6.53

11.31

2.41

0.45

0.023

2839

351,122

0.184, 0.0943

223,213

50.74

4.29

12.31

12.35

4.96

8.95

2.96

0.97

0.087

2877

321, 313

0.448, 0.-1-38

LOO, 101

# For fundamental OH stretch at 3550 em-I.* Collected in LOO ft. of water off Kalapana lava delta on 3/11/89 by G. Tribble.*:~ From Sandia drill core 150 ft deep in Kilauea Iki lava lake, 1978; obtained from D. Thomas.

\0IV

93The Sandia drill core sample was erupted at approximately 1,234 m above sea level

from Kilauea-Iki lava lake where it was air cooled, whereas the Kalapana sample was

erupted from the Kupaianaha vent (709 m above sea level) flowed down (l0.5 km) to

sea level, and then into the ocean off the Kalapana lava delta where it was quenched in

H20. Therefore, the only major difference between these two, aside from differences in

major element chemistry, is the quenching rate and the smaller H20 content in the Kalapana

sample. The lower H20 content probably reflects the longer time for degassing as the lava

moved down through a lava tube and "river" from the summit to the ocean.

Two interesting questions can be asked from observations made in this study:

(1) Why is e for the Kalapana sample (basically a degassed basalt) the same as that for the

tektite and obsidian (both high in silica) ?

(2) Why are the e's for the 2 Hawaii samples so much different?

To answer the first question the parameters that are used in the determination of the

e value were considered for the Kalapana sample. Since the density C± 50 gIL) and

thickness C±2 urn) can be measured with a high degree of confidence these two parameters

could not be expected to lead to an eroneous e value. It was possible to measure the

absorbance very accurately for this low water content glass by using a thick sample (-700

urn). The water value obtained by HTMS was therefore suspected as the source of error.

To test this hypothesis, the obsidian sample was rerun by using twice as much sample (­

150 mg). A value almost identical to the previous value ( for -75 mg) was obtained.

To see if water adsorption was responsible for the unexpected high absorbance

value, the Kilauea lki sample was left in an oven set to -120 °c for a couple of hours and

then cooled in a dessicator. When the infrared spectrum was retaken the absorbance

decreased slightly but not sufficiently to significantlydecrease the extinction coefficient

As discussed previously it is conceivable that the variation in e-values for the 3550 cm- 1

band reflects a competition among tetrahedral (Si,AI) and non-tetrahedral cations (e.g., Na)

94for hydroxyl in the melt. For example, if all OH groups are associated with silicon then on

vacuum pyrolysis the glass will produce one mole of water for every nYQ moles of OH,

viz. ,

2 (Si-OH) = Si-O-Si + H20 (ZZ)

or, more generally,

20H(melt)= 0 0 + HZO (Z3)

where 0 0 is a bridging oxygen. Therefore, on a weight basis, 34 parts of OH in the glass

will yield 18 parts of HZO. If on the other hand the OH groups are associated with Na,

then on degassing the glass one mole of HZO will be produced for evey one mole of OH,

viz., Na(OH) + H+ = HZO + Na+ (Z4)

where a Na + is exchanged for a H+ with the destruction of the Na(OH) ion pair. The

important point here is that pyrolysis methods used to extract water from glasses eventually

produce molecular H20 which is then measured by some analytical technique (e.g., mass

spectrometric) whereas infrared methods measure the concentration of X-OH groups

(where X = cation). Depending upon the relative importance of the above two mechanisms

in melts of different composition it would be reasonable to expect to find variations in the

£3550 em-I values if calibration techniques involve pyrolysis for the determination of total

water abundances. The high silica contents found in the tektite and obsidian should favor

mechanism (Z2) and yield higher extinction coefficients for the 3550 em -1 stretch. This is

indeed what we observe (£ =200-205 l/mol-cm). However, it still does not explain why

one of the basaltic glasses (Kalapana) with much lower total silica than either the tektite or

obsidian has a similar e-value,

The difference in the E values for two glasses having almost identical compositions

might be attributed to differences in the quenching rates. The quenching rate can control the

95relative amount of molecular-to-total water in the glass (see Appendix C). The Sandia drill

core which was cooled much slower compared to the Kalapana sample would be expected

to have a higher ratio of molecular-to-total water, whereas the water quenched Kalapana

sample may be expected to have a lower molecular water content. Since the total water

content in these samples is very low, infrared spectroscopy is unfortunately unable to detect

the relative amounts of molecular water in these two samples. Consequently this second

question also remains unanswered.

2. Water rich glasses

Generally, most of the glasses studied in our laboratory have less than 1.00 wt, %

total water. However, a set of glasses having at least 2 wt.% H20 have also been analyzed

(Muenow et. aI., 1980 and 1990). Due to their high water contents, it was not possible to

use these glasses in the main study. However, they provided an opportunity to determine

the integral extinction coefficients (e*) (Newman et. al., 1986) for the near-ir bands at

-4500 cm- 1 and 5200 cm-1 due to hydroxyl and molecular water, respectively. (see

Experimental Methods section: Infrared).

Two of these glasses were from the Troodos ophiolite, Cyprus (Muenow et. al.,

1990) and the third from the Scotia Sea, a back arc basin in the South Atlantic (Muenow et,

aI., 1980). They had previously been analyzed for water by high temperature mass

spectrometry and major elements using an electron microprobe (Table 10). As previously

stated these glasses were analyzed using FrIR spectroscopy; glass densities were measured

using a pycnometer. Figure 32 shows a typical spectrum (TR86-9) in the near-infrared.

Since the bands are weak it was possible to measure the areas with a higher degree of

confidence relative to measuring their peak heights. However, there is an error associated

with the background correction which is necessary here to measure the area. Table 10 also

gives the thickness, density and infrared absorbance for the 4500 cm- I and 5200 cm- I

bands. Since, we assume that the total water in the glass is the sum of the hydroxyl and

molecular water present, we can define an equation for each glass as follows:

Q.)Q

~~

..0$.0ol7.l

..0~

.6

5000Wayenumbcrs (cm.-l)

Figure 32 Near-infrared spectrum of Troodos (TR86-9) glass, which contains 2.30 wt.%H20, in the 4000-6000 cm- l region.

41000

\C0-

T.1ble 10. MajC!r Elements, volatiles and infrared data for water rich glasses

SAMPLE TR86-9 TR316 024-4

* * **

Si02 54.26 51.86 53.80

Ti02 0.54 0.66 0.60

AI203 15.76 16.35 ]4.65

FeO (Total) 7.78 7.78 8.68

MgO 6.05 6.96 8.58

CaO ]0.50 ] 1.99 10.67

NU20 1.84 2.00 1.9 ]

K20 0.22 0.16 0.19

H2O 2.30 2.11 2.04

Thicknessturn) 252 227 277

Absorbance

Area 4500cm- 1 5.81 4.83 6.76

Area 52(Xkm- 1 9.61 7.44 4.58

Density (gIL) 2758 2758 2774

*' (from Muenow et. al., 1990)

** (from Muenow ct. al., 1980)

97

98

Total water = abs(450m + abs(520m 18.02*106 (25)

(wt.% H20 ) £*4500 £*5200 d-p

Three equations were used with two unknowns (the e-values). They were fitted to a linear

reggression line on a HP-IIX calculator. This gave £*(4500) = 89.21 mol- lcm-2 and

E*(5200) =397.2 I mol- 1cm-2 with a correlation coefficient of 0.999. These values were

utilized to calculate the absolute amounts of molecular water and hydroxyl groups present

in each glass. The excellent aggreement between the total water values obtained by infrared

and from HTMS values is of course to be expected since the same water values were

initially used to calculate the integral extinction coefficients (e*4500 and £*5200)' The

amount of water present as molecular water can be used to calculate the extinction

coefficient for the bending mode at -1635 cm-l since this band was also observed in the

spectra. Table 11 gives the average extinction coefficients obtained for the three samples

and compares them with values obtained by Silver (1989) for synthetic albite glasses and

natural obsidians from the Mono Craters by Newman et al, (1986 and 1988).

The differences observed in the E and E* values obtained in this study compared to

those by Silver (1988) and Newman et. aI., (1986) can be attributed to a number of factors.

For instance, the composition of the glasses studied here is very different to previous

studies. The range in total water content was, also, markedly different (rhyolite 0.076-2.64

wt. % and albite 0.10-6.00 wt. %). The difference in quenching rate for these samples

compared to that for the obsidians and synthetic albite glasses could also have an effect on

these values.The number of samples utilized for this study was much smaller (3 samples)

compared to previous studies (-20 samples). In order to obtain more reliable data a larger

sample set is necessary.

Table 11. Comparison of extinction coefficient values obtained in this studywith those of other workers

E* (1635) E (1635) E* (4500) E*(5200)

Present Study 3,674 52 89 397(Basaltic)

Silver (1988) 3,081 49 219 268(Albite)

Newman et. al.,(1986) 2,640 55 341 248(Rhyolite)

E* --- Integral extinction coefficient (l/mol-cm-Z)

E ---- Extinction coefficient (I/mol-em)

99

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108PART2

CHAPTER IV

SPECTROSCOPIC STUDY OF MICAS AT AMBIENT CONDmONS

A. Structure and composition

Micas, which are phyllosilicates, are one of the primary storage sites of water in the

upper mantle. Unlike fluid inclusions, common to all mineral assemblages, hydroxyl

groups have structurally defined sites that are required by the mica's stoichiometery. The

word phyllosilicate is derived from the Greek word Phyllon meaning leaf, since most of its

members have a platy habit and one prominent cleavage. Some of the properties which

have made them of both practical and scientific importance for centuries are their

transparency in the visible region, low thermal and electrical conductivity (making them

excellent thermal and electrical insulators) and their perfect basal cleavage.

In simplest terms, the crystal structure of mica consists of negatively charged 2:1

layers that are compensated and bonded together by large, positively charged, interlayer

cations. A 2: 1 layer contains two tetrahedral sheets and one octahedral sheet. Each

tetrahedral sheet is of composition T205 (T= tetrahedral cation) and within each sheet

individual tetrahedra are linked with neighboring tetrahedra by sharing three comers each

(the basal oxygens) to form an indefinitely extended hexagonal sheet of Si04 tetrahedra of

the sort illustrated in Figure 33a. The fourth tetrahedral comer (the apical oxygen) points in

a direction normal to the sheet and at the same time forms part of an immediately adjacent

octahedral sheet in which individual octahedra are linked laterally by sharing octahedral

edges. The basic structural feature of mica is thus a composite sheet in which a layer of

octahedrally coordinated cations is sandwiched between two identical layers of linked

(Si,AI)04 tetrahedra (Figure 33b). The net negative charge created by the exchange of one

Si with Al is charge balanced by an interlayer cation X (K, Na, etc. in 12-fold

coordination). The hydroxyls are situated at the centres of hexagons formed by the

tetrahedra vertices (Figure 33c).

Within each 2: 1 layer the upper tetrahedral sheet has to be staggered by a distance

a/3 relative to the lower tetrahedral sheet in order to provide octahedral coordination around

o Basaloxygen:

@ ~si,:wmth •,'1,\'!;'O'TIacot'l' it

• (Si,AI)o Uxy~m

109

o Basalolygms

:!: (Si,AI)wit.osvgenbelot» It

rrrrTr1• (Si,AI)o O:rygm

Figure 33. (a ) Mica structure. (i) Plan of tetrahedrul laycr (Si,AI)401O with tetrahedra

pointing upwards and (ii) downwards and end view of layer looking along y uxis.tafterDeer al. al., 1982).

pI

o O"tlht',lrl.1J!y coordinated catious ; ntaiulv ,\ \g ..\1 or Feo ..IJJiti,lIhzl hydroxyl ions

o X ions bcloz» bottom layer (K, Na , CJ)C .\ t,';IS abat-e IIpper laver (K, Na, CJ)TII/t'k lincs :- bottom (Si I ;\1)2 Os layerThin lines : - IIfpa ( Si, AI) 2o, Ia.yer

Figure 33. (b) Plan of (i) and (ii) superimposed and linked by a layer of cations.Iafter Deeral. al., 1982).

6'\\b

4(0) 2(011)

IN"'rnbtr if iOTl$p". unit cell

o 2K6(0)

4(Sa.AI)

e

o

~\.:)

cp 0 9,o

~

~ r . r U 4(0) l(OH)

I _ 1/ '_ . -+---- __ 4(SI,:\1)~~ --.;"r-- U 0 ~ o 6 CO)

o (!}--- 0 0 0 0 lK

q ~ ~. ~_--~_O 6(0,>r I ITT i i 4o(S. ,AI)

o

Figure 33. (c) Elevation of (i) and (ii) superimposed and linked by a planeof octahedrallycoordinated cations. Composite layers are shown linked by potassium ions, and thesimplest unit cell is outlined. View is along y axis.(after Deer al. al., 1981).

--o

111the medium-sized cations. If the direction of stagger differs from layer to layer in some

regular fashion then the periodicity along the Z direction (perpendicular to the cleavage

plane) will be some multiple of 10 A and the mica is said to exhibit polytypism. Layer

structures of essentially similar composition that differ only in the layer sequences are

called polytypes.

The general formula which describes the chemical composition of micas is:

X y 2-3 Z 4 °10 (OH,F,CI)2

where X is mainly K or Na (interlayer cation)

Y is mainly AI, Mg, or Fe (octahedral layer)

and Z is mainly Si or Al (tetrahedral layer).

These micas can be further subdivided into dioctahedral and trioctahedral classes in which

the number of Y ions is 2 and 3, respectively. If all three octahedra are occupied, i.e., have

octahedral cations at their centres, the sheet is classified as trioctahedral. Phlogopites,

which are trioctahedral, have the ideal formula: KMg3AISi301O(OHh. If only two

octahedra are occupied and the third is vacant, the sheet is classified as dioctahedral.

Muscovites, which are dioctahedral, have the ideal formula: KAI2AISi301O(OH)2 (Figure

34). Biotites, which are also trioctahedral, have the general formula

K(Mg,Fe)3(AI,Fe)Si301O(OH,F,CI)2' In practice most micas do not contain exactly 3 or

2 octahedral cations, but 2.5 forms a convenient boundary between the subgroups because

few homogeneous micas have octahedral cation totals near that value.

B. Study of micas

The two principle mica groups, dioctahedral and trioctahedral have been relatively

easy to identify, with an optical microscope. For instance, the angle of the optic axes ranges

from 500 to 750 for the former group, whereas values below 150 are observed for the latter

group. Previously, apart from elemental analysis, density, hardness, and thermal stability

were some of the parameters used to differentiate these two groups. Mass magnetic

susceptibilities were measured for thirty-eight micas by Hood and Custer, (1967). They

found it to be a function of the oxidation state of the iron, with ferrous iron imparting a

slightly higher susceptibility than ferric iron. One would like to have a better understanding

~IDA

11

o

o

o2l<'+

2k+ 0

,I,3 ~i4+ , ,.H~',

U I Jrll I \

60'- -' ~

lX'

3Ji", AI" 3Si~'JAt~

60'· 60"

40~?Olr

40~°i!OH- -1 All.

-lAl"

4. O~-20/l-

60'-

J.5i 4' , AI"

,

~;, {(~'·~OI.~ ......

.- .,--.............. ,,;) o-a"--coo. 'IS --'-

Figure 34. The idealized structure of muscovite: Repeat unit on c is two tetrahedral layersand one octahedral (dioctahedral) layer with interlayer cations (-lOA). (after Phillips andGriffen, 1981).

hJ

113of the structure of these micas and how the chemical composition affects the structure.

Skow (1960) stressed this need when he said that "such knowledge would lead to a

practical, reproducible, scientific test to evaluate sheet mica for specific end uses".

The basic structure of mica was obtained by Pauling (1930). Accurate structure

determination didn't become possible until Radoslovick, (1960) and Steinfink, (1962)

published their work. Recent x-ray crystallography work on the elucidation of the structure

of micas has been published by Hazen and Finger (1978) and Schroeder (1990), for

phlogopite and muscovite, respectively. Vedder and Mcdonald (1963), Novak (1974),

Bailey (1984) and Smith et. al. (1987) give a good summary of the relation between

composition and structure of micas.

Some of the other techniques utilized to study micas include Mossbauer, electron

spin resonance, nuclear magnetic resonance (Sanz and Stone, 1983) and infrared and

Raman spectroscopies (Rossman, 1984). Infrared spectroscopy has been used extensively

to investigate micas, but only a handful of Raman studies have been published (Loh, 1973;

Haley et al., 1982; Robert and Kodama, 1988 and Tlili et. al., 1989). Rossman (1984)

wrote the first review article on the application of vibrational spectroscopy to study micas.

Farmer (1974) had previously discussed the application of infrared spectroscopy to layer

silicates as a group. Aines and Rossman (1984) have discussed the infrared spectra ofOH

and H20 in minerals. Nakamoto (1986) gives a general discussion of the vibrational

spectra of water molecules and the OH ion.

C. Infrared spectroscopy

The isolated OH- ion has a single motion, O-H stretch, which absorbs infrared

radiation in the -3700 cm- l region. The exact position of this band is a good indicator of

the environment of the hydrogen bond because in crystalline solids the H atom is usually

hydrogen bonded to another ion usually oxide. So the exact frequency of the O-H stretch

will depend upon the strength of the hydrogen bond. In general, the frequency of the OH

stretch occurs at lower energies in systems in which the OH ion is more strongly hydrogen

bonded. Nakamoto et. al., (1955) and Novak (1974) have presented the correlation

between the frequency of the O-H stretch and the strength of the hydrogen bond as

114measured by the O-H---O distance. Furthermore, if the OH groups are structurally

orientated within the host, the amount of incident radiation absorbed can be strongly

dependent upon the relative orientation of the OH dipole and the direction of linear

polarization of the interrogating light.

Previously, the stoichiometric hydrous components were studied by conventional

methods of infrared spectroscopy. These studies used the KBr pellet method to determine

the specific species (OH of H20) present, and to use the OH stretching modes as probes of

local cation ordering. However, single crystal studies are preferred because they generally

avoid the problem of adsorbed water on the fine powders used with the KBr method. Also

they provide information about the orientation of the OH in polarized light, and at known

path-length, they provide better quantification of the hydroxyl content

Much attention has been given to the analyses of OH-stretching bands of micas and

other phyllosilicates because it appears that each of the various bands in the OH region

(3800-3200 em -1) can be attributed to a specific atomic grouping around the OH site

(Velde, 1983 and Robert and Kodama, 1988). The stretching frequency of the OH groups

occurs in an isolated region of the spectrum facilitatingidentification and analyses.

Infrared spectroscopy has been used to identify individual micas, estimate chemical

composition, determine ordering and orientation of small molecule-like units in a mica

(Vedder, 1964, Jorgensen, 1966, Robert, 1976, Tlili et. al., 1987 and Robert and

Kodama, 1988). Hydroxyl stretching frequencies are affected by the octahedral cations to

which the hydroxyl groups are coordinated (Vedder and Mcdonald 1963, Wilkins 1967,

Gilkes et. aI., 1972 and Slonimskaya et. aI., 1986), by interlayer cations in the micas

(Chaussidon, 1970 and Smith et. aI., 1987), and by the configuration of, and charge

distribution in the surrounding tetrahedral lattice (Stubican and Rustum 1961, Farmer and

Velde 1973 and Tateyama et. aI., 1976). The first two factors are, predominant in the

trioctahedral series. Farmer (1974) summarizes the spectra in this region. In the spectra of

trioctahedral micas, the primary OH stretching frequency is at -3714 em-I, but can vary

with substitutions in the octahedral and tetrahedral sheets. Minor components which can

115

occur as three broad bands in the 3450 to 3620 cm- l region are a result of associations of

OR with vacancies in the octahedral sheet, especially in high Fe-content biotites.

Several studies have focused on the systematic study of IR spectra, of a series of

synthetic micas (Langer et. al., 1981, Robert et. al., 1983 and Liu et. al., 1987). Velde

(1980) has correlated all dimensions with the IR spectra of synthetic dioctahedral

muscovite-phengite, margarite, and paragonite. The spectra indicated that changes in the b

axis are due to ordering in both the tetrahedral and octahedral sheets.

Velde (1983) examined the crystal chemical factors which influence the energy at

which a particular OR stretch occurs. The average electronegativity of the group of

octahedral ions bound to the OR gave the best correlation with the OR stretch frequency in

di- and tri-octahedral potassic micas. A frequency shift of 96 cm- l per electronegativity unit

with trioctahedral micas and 170 cm- l for dioctahedral micas were observed. Lesser

effects were observed for substitution in the tetrahedral sheet where changes in unit cell size

were more important than electronegativity effects.

The orientation of the OR groups in micas have been determined from the infrared

spectra (Serratosa and Bradley 1958, Bassett 1960 and Giese 1979). Phlogopite provides

a simple demonstration because in the ideal end-member composition each OH occurs in

just one site associated with three Mg ions. Serratosa and Bradley (1958) showed that there

is substantially no absorption when the electric vector of the incident light is parallel to the

cleavage plane, but that absorption increased as the mica flake is tilted such that a

component of the electric vector is normal to the cleavage (Figure 35). Because the OR

group is excited into vibration only when the electric vector is aligned with the O-R bond,

this result demonstrates that in phlogopite, the OR group is oriented essentially normal to

the cleavage plane. In phlogopite the band is weak at 3714 cm- l and increases as its

cleavage plane tilts toward the infrared beam (Vedder 1964, Vedder and McDonald 1963).

Comparable behavior was observed for other trioctahedral minerals. Vedder (1964)

redetermined orientation of OH groups, and concluded that there was a 100 angle between

the OR unit and optic c* axis. The direction perpendicular to the a and b axes. It deviates

116

PHlOGOPITE

II I ...." I,." .

MlISC:OVITF.100'..------------------,

0"30"C!l"fi(I"

«' r..lm "ur.knell ... 004 1ft"'

,. ..... ,;.,. ...... ~..., "

" \

r '1 I

I,,,,"t

UIO

10

m:

pJVV,,IIIII,

,.,.

60

40 - 0---. "!Ie

1.0

,v

IU' ,

~1\i'O

coooI I ,

.",.1

WI\VENUMBERS (eM-I)

Figure 35. Infrared transmission spectra in the OH stretching region of cleavage plates ofphlogopite and muscovite which have been rotated at angles indicated to show the stronganisotropy of the OH stretching mode in phlogopite. (after Scrratosa and Bradley, 19S5).

117from the crystallographic c axis by a few degrees. The spectrum of the dioctahed.ral micas,

for instance muscovite, indicates that the OH groups are inclined (-780 ) from the normal to

the cleavage plane because of the stronger intensity of the OH band due to vacancies in the

octahedral sheet. The dioctahedral OH frequencies are also typically lower because of

stronger hydrogen bonding. Muscovite has its OH stretching frequency at uO-H = 3625

em-I. Loh (1973) explained the difference in OH-stretching frequency from uO-H =3714

em-I in phlogopite to uO-H =3625 cm- 1 for muscovite as a consequence of the weak

bonding between the tilted hydrogen and the two nearest nonbridged oxygens for

phlogopite. In accord with the chemical analysis of many biotites, which indicated cation

occupancy intermediate between di- and tri-octahedral, the infrared spectra show two types

of OH absorptions, one at higher energies indicating trioctahedral orientation almost

perpendicular to the cleavage plane and the other at lower energies indicating muscovite

orientation almost parallel to the cleavage plane (Serratosa and Bradley, 1958 and Wilkins

1967, Figure 36). Since the strength of O-H absorption depends on the orientation of the

OH relative to the electric vector of the radiation, the orientation distribution of the OH

groups must be taken into account in comparing the strengths of absorption in materials in

which the hydroxyl is incorporated in different ways.

In end-member trioctahedral micas, a single OH stretching band at about 3714

em-I results from a single OH environment, Mg3(OH). In chemically more complex.

micas, multiple bands arise from multiple groupings such as Mg2Fe(OH) and MgFe2(OH)

(Wilkins and Ito, 1967). There have been extensive rr;·~a OH spectral data presented which

have been previously reviewed by Farmer (1974). The use of the OH

bonds in the 3800-3200 cm- l range to establish cation site occupancy has assumed that the

intensities of the individual OH absorptions are proportional to the concentration of the OH

groups in each chemical environment. When multiple cation-OH groupings are involved, a

number of studies have indicated that, in fact, there are different absorption intensities

118

100..--------------------,

BIOTITE

30

50

40

--- _....... ,I /-----

70 _ I I- - \ I I

I

\. I II' I

I' I'I \ ,il- .. I TI 13580

I rI I\ I

1/t

3700

60

Film ,,."cknU5 - 0 10 ",'"

zo

to

4000 3500 3000 em·'

Figure 36. Infrared transmission spectrum in the OH stretching region of a cleavage plateof biotite which has been rotated at angles indicated. In accord with the chemical analysesof many biotites which have cation occupancy intermediate between di- and tri-octahedral,the infrared spectra shows two types of an OH absorption, one at higher energies indicatingorientation nearly normal to the cleavage, and one at lower energies indicating orientationinclined to the cleavage. (after Serratosa and Bradley, 195R).

119associated with each particular cation-OR grouping (Rouxhet, 1970; Gilkes et aI., 1972;

Sanz and Stone, 1983).

Studies of near- and mid-IR mica absorption spectra have provided insight into the

causes of vibrational modes of OR aluminosilicate tetrahedra, and metal octahedra which

helped to constrain the locations and orientations of these intrasheet components (e.g.

Fanner, 1974). In addition to these studies, with the advent of the Fourier transform

infrared spectrometer has made possible routine measurement of absorption spectra in the

far-ir (50-300 cm- 1) region (Velde and Couty, 1985 and Schroeder, 1990). The

instrumentation and experimental methods used in the spectral region below about 400

em-I are different consequently, the far-infrared spectra are usually considered separately.

Comparatively few mineralogical studies have been performed in this range which

responds to vibrational motions of heavy ions and the motions of larger units of the

structure. Larson et. aI., (1972) surveyed the far-infrared (FIR) spectra of three micas and

found that they were distinct. Farmer and Velde (1973) determined that there is tetrahedral

AI, Si ordering in margarite based on the sharpness of the IR spectra and the lack of AI-O­

Al vibrations. In fact most of the current infrared work on micas includes a discussion of

ordering. Infrared spectra are also sensitive to mica polytypes. For example, Velde (1980)

observed that the spectra of synthetic margarites vary with the polytype, and sensitive

bands occur at 446 and 612 cm- 1 in the margarite spectra and at 643, 737, and 802 cm-1 in

the muscovite.The positions of low energy vibrations such as those found in these studies

are needed for input parameters into thermodynamic model calculations such as those of

Keiffer (1982).

D. Raman spectroscopy

Loh (1973) was the first person to utilize Raman spectrometry to study micas. He

used a combination of Raman and infrared spectroscopy to study the vibrational spectra of

micas and other phyllosilicates. Since most of the previous work had involved the infrared

region 300-4000 em -1, Loh (1973) studied the far infrared spectra of micas. His data

120suggested that the vibrations in micas can be described as molecular vibrations of the

following type:

(i) The distorted octahedron M06 with S6 symmetry below 210 em-I; where M is the

octahedral cation such as Mg2+, Fe2+ or Al3+.

(ii) Isosceles triangle O-H-O in dioctahedral micas containing octahedral Al3+, for example

muscovite, between 200-300 em-I.

(iii) Distorted tetrahedron SiO4, C3v synunetry and OH ion from 300-4000 em-I.

The first systematic study of natural micas by Raman spectrometry was carried by

TIili et. al., (1989). This study showed that micas can indeed be throughly respectable

minerals for Raman spectroscopic work (Smith et, al., 1987 and TIili et. al., 1987). The

micas studied were dominated by (Mg, Fe, AI) and (Si, AI) in the octahedral and tetrahedral

sites, respectively. Raman spectra in the ranges of 50-1250 cm- l and 3500-3750 cm- 1

were presented, with attention focused on the differences between dioctahedral and

trioctahedral micas and on the apparent trends of variation of each major Raman band with

composition.

121CHAPTER V

EXPERIMENTAL METHODS

A. Infrared spectroscopy

This technique was previously discussed in detail in Part 1 of this dissertation. See

pages (18 ) - ( 25).

B. Raman spectroscopy

In Part 1 infrared spectroscopy was discussed where infrared radiation, of a

spectrum of energies, interacts with matter by being absorbed in order that the molecule can

undergo a particular stretching or bending motion. The infrared absorption band gives

information about the vibrational energy levels of the molecule. Similarly, Raman

spectroscopy can also give us information of vibrational and rotational energy levels,

although the experiment is very different. Raman spectroscopy uses a monochromatic

visible light source, a laser, that undergoes scattering when it interacts with the specimen.

This scattering results from the interaction of the electric vector of the electromagnetic wave

with the induced dipole of the species being observed. Vibrations of atoms in a molecule

induces oscillating electric dipoles. Such oscillating dipoles become new sources for

emitting radiation, that is, the scattered light. Three types of scattering are possible (Figure

37):

(1) Elastic or Rayleigh scattering. Same frequency (wavelength) as the incident light.

(2) Inelastic or Raman scattering. Lower frequency (longer wavelength) than that of the

incident light. (Stokes-Raman scattering).

(3) Inelastic or Raman scattering. Higher frequency (shorter wavelength) than that of the

incident light. (anti-Stokes-Raman scattering).

Although Rayleigh scattering occurs more effeciently, its the inelastic scattering that

contains information on the vibrational and rotational states of molecules. So, even though

the signal is weak, through the use of an intense monochromatic laser source in conjuction

with a very sensitive detector, spectroscopists have been able to obtain the Raman spectra

of various polyatomic species. The size of the particle or molecule illuminated, location,

frequency and intensity of the incident light are some of the variables which control the

intensity of the scattered light.

122

Figure 37. Diagrammatic presentation of three types of li~ht scattering: anti-Stokes Ramanscattering, Rayleigh scattering and Stokes Raman scattenng (after Tu, 1982).

--··------1 El~ctron'c__. . . .._=~ (ICII'!'" Slf]l~

\

'/lbrfltlf)nQl Quantum N..".,b@,"",thIn rh~ EI!ClronlC

Ground S'OI~

. I I.r--; Vir °JO

.r;,0 Slol~

r-,. Or"'..

-- - - - -- -- - - - -- .-- I-I.-- --

Slo"es Anlo ° SIO"~'

Lln@ Line

Romon Scottermq

Figure 38. Energy diagram of a molecule showing the origin of Stokes and anti-StokesRaman scattering. The molecule is momentarily elevated to a higher energy (virtualstatelbut it never reaches an electronic excited state (after Tu, 1982).

1231. Raman scattering

The Raman scattering effect arises from the interaction of the incident light with the

molecular species in the illuminated matter. In nonresonance (laser line has wavelength

different from the absorption edge of the analyte) Raman scattering, the energy of the

incident light is not sufficient to excite the molecule to a higher electronic level. Instead,

Raman scattering results in changing the molecule from its initial vibrational state to a

different state (Figure 38).

2. Stokes and anti-Stokes lines

In order for a molecule to exhibit the Raman effect, the incident light must interact

with the induced dipole or produce a change in molecular po1arizability. Consider carbon

dioxide as an example. The polarizability can be quantitatively visualized as a change in the

shape of the electron cloud. As shown in Figure 39, the electron cloud around the C02

molecule is alternately stretched or shrunk inphase upon interacting with the oscillating

wave of the electric component of the electromagnetic wave. This results in scattered

photons with altered frequency. The scattered light contains a small portion of light due to

Raman scattering in addition to that due to normal Rayleigh scattering (same frequency as

the incident light). The Raman scattering spectrum contains both the Stokes and anti-Stokes

lines; their frequencies correspond to the sum and the difference of the incident light and the

allowed molecular vibrational frequencies. When photons interact with a molecule, some of

their energy can be converted into various modes of vibration of the molecule. As seen in

Figure 38, the scattered light loses energy equivalent to the energy given to excite the

molecule to various vibrational energy levels (Stokes Raman effect). If the energy is

transferred to the incident photon from a molecule in the excited state, then the scattered

photon has more energy than the original incident photon (anti-Stokes Raman effect). This

requires that the scattering molecule already be in an excited state. For typical applications

this is seldom the case; therefore, Stokes Raman scattering is much stronger than anti­

Stokes Raman scattering.

124

~..'(J ....: I .':.,

hv ;- .,;

A/\~;"c\J V V~·· .~

~ \ f..~ •. 'T·"r '.#'0··

U

Figure 39. An example of change in polarizability. The Raman effect requires a change inpolarizability (change in electron cloud) when illuminated with light during a vibrationalnormal mode (after Tu, 1982).

fl~

Stokes Line

RayleiQhLine

6CmI 500 0 500

cm' 19.992 20.492 20,992

Figure 40. The Raman shift (frequency difference) for Stokes and anti-Stokes lines arealways symmetrical (after Ttl, 1982).

125Only a small fraction of photons are scattered inelastically, so Raman lines are

usually very weak (only 10-6 of the intensity of the Rayleigh line). The majority of the

scattered light is the same as the original incident light in terms of photon energy. This is

the reason why a laser is used because it is an intense and monochromatic source of light,

which has high photon density.

The Stokes and anti-Stokes lines involve the same vibrational-energy difference

(Figure 38). Therefore, the difference between the incident-light frequency and the

scattered frequency in the Stokes and anti-Stokes Raman scattering is identical. That is, the

difference in frequency for the Stokes and anti-Stokes Raman scattering is symmetrical

with respect to the exciting line. For instance, a compound illuminated with the blue line of

an argon ion laser 488 nm (20,492 cm-1) produces two Raman lines at 19,992 cm-1 and

20,992 cm -1. Although one cannot see the significant correlation between these lines

merely by inspecting the two values, one can readily see an interesting result by taking the

frequency difference between the scattered and the incident light (Figure 40). Normally,

Raman scattering is expressed by this wave number difference (L\ wavenumber or Raman

shift) rather than the absolute wave number. The Raman spectrometer automatically gives

the wavenumber shifts (or Raman shift), which contains information about the vibrational

level of the molecule.

The frequency of the Raman line is independent of the wavelength of the incident

light. One should not confuse this with the fact that the intensity of Raman scattering

depends on the wavelength (A.) of the incident light and, actually, is inversely proportional

to 11.4. Thus whether one excites a molecule with green light (514.5 nm or 19,436 cm- 1) or

blue light (488.0 nm or 20,492 cm- 1), one will obtain a Raman shift line at exactly the

same wavenumber difference. This can be readily understood from Figure 41.

From Figure 37, one sees that both Stokes and anti-Stokes Raman lines have

identical frequency values. However, the Stokes line has a higher intensity than the anti­

Stokes line at room temperature. The Stokes line originates when a molecule at a low

126

--7""\-- -

r----r<--Energy

ofBlue light

- ---- n F~ -

-r---Energy

01Green light

J__

Figure 41. The frequency of the Raman-scattered light is independent of the excitationwavelength (after TlI, )lJR2).

127

_... _-_ ..... -.--r"-,..".,- .......... -.- .. - ... -. .~ "'to?..) ~. . :....L- - • -yoZ A _ T.

A --1-, V. V· I-.L:-6"...----''---Yo, .........-1_"'IJ.:...' V-O .....:;:.:-oox-_----:'--_ yoO

Stolles Effect Anti-Stolles Effect

AI ~ '..."e'atu'! - The papulation of

level V'O i. mo,e than 'hat 01 Vol

Iftolec:uln at ene,o.,

~ hiSlher temper~ - The population of molec:ulu at ene"n

'eve' Vol ,ne'.atel ,elo'i.e '0 that 0' -UlY VoO

A HiOhe,j ~nten~'IY

Figure 42. The origin of Stokes and anti-Stokes effects and their intensity difference. Therelative intensity of Stokes and anti-Stokes lines vary depending on the temperature. (afterTtl, 19X2).

128

vibrational energy is elevated to higher hUl (~u in Figure 42) by interacting with the

incident light, whose energy is equal to huo' On the other hand, the molecule at a higher

vibrational-energy level gives away the energy hUl; therefore, the molecule becomes lower

in vibrational energy, and the scattered light increases in energy by hUl' At low

temperatures (or at room temperature), more of the molecules are at lower vibrational­

energy levels than at higher vibrational-energy levels (Figure 38). Thus a larger fraction of

molecules will have Stokes-type transitions than anti-Stokes transitions, and the Stokes line

will have higher intensity than the anti-Stokes line. This can also be seen by examining the

Boltzman distribution law, which states that the relative population of molecules with

higher energy increases as the temperature is increased:

N/No = e -~ElkT (26)

where Ni is the number of molecules at energy state Ei and No is the number of molecules

at energy state Eo ; NiNo is the fraction of molecules at the energy state Ei; k is the

Boltzmann constant; T is temperature in the Kelvin scale; and ~E is the energy difference

between energy states Ei and Eo.

Therefore, the intensity of anti-Stokes-lines increases (or the Stokes-line intensity

decreases) as the temperature is elevated. The ratio of the anti-Stokes to the Stokes line is

directly related to the fraction of molecules at higher vibrational-energy levels.

3. Classical mechanics of Raman scattering

Classical mechanics will be utilized here to explain the scattering of light by a

molecule. This treatment is relatively easy and both elastic and inelastic scattering effects

(Rayleigh and Raman scattering) can be explained by one equation.

If a molecule interacts with light, the electric field of photons will exert oppositely

directed forces, an induce dipole caused by seperation of the centers of negative and

positive charges of the electrons and the nuclei, respectively. As a result, the polarized

molecule will have an induced dipole moment (P) caused by the vibration of atoms in the

129molecule. This induced dipole moment, in the presence of electric field E, is proportional

to the electric field E and to a property of the molecule called the polarizability a :

P =a E (27)

The electric field is an oscillating function dependent upon the frequency of the light uo:

E = Eo coszx Uo t (28)

Substituting this into (2), we obtain:

P = a Eo coszn Uo t (29)

The polarizability a is dependent upon the position of the nuclei in the molecule. For a

molecule containing N atoms, there are 3N degrees of freedom available to the nuclei. Of

these, 3N-6 (3N-5 for a linear molecule) result in vibrations of the molecule. In general,

the vibrational motion of all but the simplest molecules (e.g., diatomic) is quite

complicated. However, by group theory it is possible to reduce this complicated vibrational

motion to a set of independent normal modes of vibration.

Instantaneous positions of the nuclei can therefore be expressed relative to their

equilibrium positions in terms of the normal coordinates Qi' where i=I,2, .... ,3N-6.

Considering a diatomic molecule with the single normal coordinate Q 1, the dependency of

<X on Ql is expressed as a series expansion:

a = ao +(8a /8Ql)0 Ql + (30)

where ao is the equilibrium value of the polarizability.

The position of the nuclei is time dependent because the molecule is vibrating with

frequency Ul' Information on the frequency of vibration can be obtained from knowledge

of the forces between the vibrating nuclei, and the application of the classical mechanics of

small vibrations. This motion can be expressed as:

QI =Qlocos21tUI t (31)

where Ql° is the maximum vibrational amplitude. It is seen, therefore, that a also

130

oscillates at the frequency u1' Substituting equations (30) and (31) into eqn. (29), we have:

P =[ao +Coa / OQl)o Ql lEo cos21t Uo 1.. (32)

From cos S x cos c!> = 1/2 [cos(S +c!> ) + cos(S - c!>)]

P= ao Eo coszn U o t + 1/2 Eo Q10(oa / dQ1)0

x [cos21tt(uo +u1) + cos21tt(uo - u1)] (34)

This classical derivation for a diatomic molecule predicts three basic light scattering

modes due to the induced dipole moment P oscillating at frequency u1:

1. The a o term produces scattered light unshifted in frequency (Rayleigh scattering).

2. If OcxlOQ does not =0, Raman scattering occurs, and the incident light of frequency U o

is shifted to scattered light with frequency Uo

+u 1(anti-Stokes) and lower frequency Uo-u1

(Stokes).

4. Instrumentation

Broken down into its basic components, a Raman spectrometer consists of a light

source, a collection optic, a dispersing optical element and a detection system.

The driving force behind most changes and improvements in the instrumentation for

Raman spectroscopy is the inherent weakness of the scattering phenomenon. Since very

few photons are scattered inelastically, the detection system employed must be extremely

sensitive and able to detect small numbers of photons over the dark noise background.

Most attempts to improve the sensitivity have involved: increasing the effective collection of

scattered photons, decreasing the transmission losses of the spectrometer, or increasing the

sensitivity of the detection system.

5. Sources

Most applications of Raman spectroscopy have involved the use of continuous

wave (cw) gas ion lasers, either argon or krypton.

1316. Monochromators

A monochromator is required to disperse collected light across its exit slits for

sequential presentation to a detector or to disperse it across an array detector. Since the

elastically scattered component is also collected the monochromator must be able to

discriminate effectively against this large signal.

The increasing use of array detectors has forced the development of efficient, low

stray light level spectrographs. These are usually composed of a first-stage double

monochromator, followed by a third grating that acts as a "normal" monochromator and

disperses the bandpass of the first stage onto the face of the array detector. This setup is

used in the state-of-the-art optical multichannel array detection instruments (OMA's).

7. Multichannel Raman spectrometry: A triplemate spectrograph

The first two stages of this instrument act as a variable band-pass filter, and the

third stage provides the dispersion. Significant flexibility is provided by rapidly

interchangeable gratings in the triple spectrograph; with a diode array detector, for instance,

resolution and coverage range from 1.4cm-l per channel with a spectral coverage of 980

em-I to 5.6 cm- l per channel with a 3902 cm-l spectrum.

The multichannel spectrograph allows one to record a substantial segment of the

Raman spectrum (ideally, the entire Raman spectrum) simultaneously without scanning the

spectrometers wavelength setting (Campion and Woddruff, 1987, Sharma and Urmos,

1987, Schlotter et. al., 1988,and Sharma et. aI., 1988 ). It will outperform the scanning

spectrometer in the time required to obtain a spectrum with a given signal-to-noise (SIN).

Although, this factor depends upon detailed specifications of the detector and the

spectrograph, a typical improvement in spectral acquisition time is a factor of 1000.

Another way of stating the multiplex advantage is that the signal-to-noise ratio for a given

spectral acquisition time will be better in the multichannel case by a factor equal to the

square root of the number of resolution elements recorded simultaneously. This factor

which is typically 30 can be the difference between the observation of a good spectrum and

worthless noise. This result can be a crucial advantage when extreme sensitivity or sample

durability is of concern. The major drawback of such a system involves a

132resolution/bandwidth trade-off. If high bandwidth or spectral coverage is required, then

resolution must be sacrificed. Although 1024 elements are usually present, usually only

1000 are actually active. So, resolution is reduced. In spite of these problems, the gain in

sensitivity is significant. Continuing advances are being made in the sensitivity of these

array detectors. The recent introduction of charge coupled device (CCD) will improve the

results being obtained by multichannel detectors (Bilhorn et. al., 1987).

The introduction of lasers, which increased the intensity of the exciting radiation,

and better detector systems led to the widespread use of the technique in the 1970's. The

introduction of the diode array detector in the 1980's, brought the multiplex advantage back

to Raman spectroscopy that it had given up in the change from photographic to

photoelectric detection. Recently, the coupling of microscopes to Raman spectrometers

have greatly expanded the versatility of the technique.

The geometry used in experimental micro-Raman spectroscopy has facilitated the

use of special cells, such as high pressure diamond anvil cells (DAC's) and/or low/high

temperature cells for in-situ spectroscopy for minerals ( Sharma, 1989a,b).

Raman spectroscopy has been limited in its application by fluorescence. As a

phenomenon, fluorescence is approximately 106-108 times stronger than Raman scattering.

The presence of trace transition metals, coatings on polymers, additive, etc., have been

observed to fluoresce so strongly that it is impossible to observe the Raman spectrum of a

major component. The introduction of Fl'-Raman spectroscopy, which utilizes a near-ir

excitation source, has helped to overcome/reduce this problem.

C. Instrumental methods

The multichannel micro-Raman facility situated in Hawaii Institute of Geophysics

(HIG 107) was utilized for this study. The instrument is capable of measuring Raman

spectra in either 1800 or 1350 scattering geometries. The system is shown in Figure 43. It

employs a modified Leitz Ortholux-I microscope which is optically coupled to a spex

Triplemate spectrograph and an optical multichannel detector and analyzer (OMA ill, model

133

Analyzer

Measurlng-ftelddIaphragm ~~~~3

(B

Remen spectrometer

TV Monitor

Laser

Figure 43. Micro-Raman apparntus for lise in 1350 and 1ROo scattering geometries with aCW laser, and a multichannel Raman spectrograph.

I

1341421 detector and model 1460 analyzer, EG&G Princeton Applied Research). The

microscope was modified to allow sufficient working distance between the objective and

stage to accommodate either a diamond anvil cell, high temperature fumance, or liquid

nitrogen or liquid-helium cryostate.

The sample was excited with an Argon ion laser (Spectra Physics 2020) with an

average power of 30-50 mW using 1350 geometry. The incoming laser light was focussed

upon the sample and the scattering light was collected by a 18X objective. The laser line

utilized changed from 457.9 nm for the hydroxyl region (3000-4000 em-I) to 488 nm for

the low region (130-800 cm "). However, the O-H stretching region was calibrated with

the 514.5 nm (green) line. This laser line was used to obtain the spectrum of optical grade

calcite. The calibration involved shifting the position of each calcite line (154,281,711 and

1085 cm- 1) by adding 2403 em-1 to each line. This value represents the difference in

wavenumbers between the 514.5 (19,436 cm- l) and 457.9 nm (21,839 cm- 1) lines. The

calibration utilized the cubic-fit calibration of the abscissa, which uses a third-order

equation. All the work reported here was carried out with the 1200 grooves (g)/mm grating

in the spectrograph stage, with the wavenumber calibration accurate to within ± 1.5

cm- l(Sharma et. al., 1988). A diode array 1024 element detector, cooled down to -25 °C,

was utilized for most of the studies. Some of the mica samples at ambient temperature

were studied using a 2-dimensional CCD, liquid nitrogen, cooled detector.

The highly sensitive detection system utilized by this spectrometer provided an ideal

opportunity to study these samples which had been, previously, termed as not very good

samples for Raman spectroscopy, due to their poor response when illuminated with a laser.

This study provided the opportunity to utilize two different detectors. A diode array 1024

element detection system was used initially. The multichannel capability of the spectrometer

allows one to record the whole spectrum simultaneously rather than a scanning

spectrometer which scans a wavenumber at a time. Normally, a scanning spectrometer

takes approximately 20 mins to obtain I scan whereas with this system it was possible to

135obtain the whole spectrum in one minute. The signal-to-noise (SIN) ratio can be improved

greatly by taking multiple scans. Also, the detector when cooled to -25 °C meant that the

SIN level could be improved. The other detector used in this study was a CCD. This

detector has a higher sensitivity, relative to a diode array, and can be cooled down to -120

°c by using liquid nitrogen. However, it was not possible to scan for long periods of time

meaning that fewer number of scans could be taken. This resulted in signal to noise (SIN)

levels that was not as good as that obtained by the diode array detector. Charge coupled

devices (CCD's) have higher read out noise compared to diode array detectors;

consequently, noisy spectra result for spectra obtained with small exposure times. The

number of scans, exposure time, slit width and laser power varied somewhat as these

samples were scanned over numerous sessions. Despite obtaining good spectra frequently,

some technical difficulties were encountered when recording their spectra (e.g., high

background noise, high fluorescence or strong absorption). This happened especially

when the crystals were Fe2+-rich. For some phlogopites, it wasn't possible to find grains

that gave good spectra.

The infrared spectra were obtained using a Perkin Elmer PE1720X spectrometer

using the same set up and parameters eluded to in Part 1 of this dissertation. The infrared

spectra were obtained in the region, 400-7000 em-I, which covers the mid-infrared and

near-ir regions.

136CHAPTER VI

RESULTS AND DISCUSSION

A. Samples

A collection of phlogopites from nodules in South African kimberlites, and

muscovite and biotite samples were studied using a combination of infrared and Raman

spectroscopy. The phlogopite samples were previously analyzed by electron ~croprobe

and high temperature mass spectrometry (HTMS), to obtain their elemental and volatile

abundances, respectively (Matson, 1984 and Matson at. al., 1986). Table 12 gives the

chemical analysis of a representative phlogopite mica, BD3088. The muscovite and biotite

samples were obtained from the Geology and Geophysics mineralogy collection. Since

these latter two samples were not characterized before, their volatile and major element

analyses were obtained, by HTMS and electron microprobe, respectively, in this study (see

Appendix H).

The phlogopite micas had formed under several distinct modes of occurence in

kimberlites. Differences in the modes of occurrence of the micas generally reflect distinct

crystallization environments and/or post crystallization histories. These were further

reflected in their chemical compositions and the textural relationships of the micas with

surrounding minerals (Matson, 1984). The relative concentrations of species occupying the

hydroxyl site of the mica structure (i.e., OH-, F, Cl", and perhaps 0 2- ) are determined by

the relative activities of these species in the system from which the micas crystallized.

Phlogopite mica is considered to be a primary volatile bearing phase in the upper mantle

because of its common appearance in xenoliths transported to the surface by kimberlites

(Matson, 1984). Calculated on the basis of its pure endmember formula

[K2Mg6[Si6Al6~O](OH)4;Deer et. al., 1982), phlogopite mica contains roughly 4.2 wt.

% H20 .

Water structurally bound as hydroxyl ions is an important component of hydrous

minerals such as mica. Its concentration should be reported, therefore, as part of the

compositional analysis. In addition to providing total volatile content, HTMS has the

potential to provide information on the site occupancy of specific volatile species within the

137Table 12. Chemical analyses of a phlogopite mica (wt.%)

Sample 8D3088

Si02 41.48Ti02 1.36

AI203 12.61

Cr203 0.23

FcO* 4.03MnO 0.08MgO 24.8NiO 0.11C10 0.02Na20 0.13

K20 10.98F 0.37Cl 0.06H2 O 4.25

Sum 100.51

0= F,C1 0.17Total 100.34

Ionic Formulas Calculated on the Basisof 24 (0, 01-1, F, CI)

Si 5.855Alz 2.098Aly 0Ti 0.144Cr 0.026Fe 0.476Mn 0.01Mg 5.218Ni 0.012Ca 0.003Na 0.036K 1.977F 0.165CI 0.014OH 4.002

Cation Totals

Z 7.95y 5.87X 2.02on-r-ei 4.IRMg# 91.6

(from Malson ct. al., 1986)

138mineral's structure. Complete chemical analyses of hydrous minerals require the

determination of cation components contributing to the structural framework.

Compositional analyses of the micas investigated include major elements (determined by

electron microprobe) and volatile components believed to occupy structural sites

(determined by HTMS). Chemical formulas of the micas were calculated on the basis of

24-anions according to the method of Deer et. al., (1982).

Most of the spectroscopic measurements have been performed on natural micas.

There is, however, a need for the systematic study of simple and well characterized micas.

In this study, attempts have been made to obtain the vibrational spectra of the phlogopites,

muscovite and biotite samples eluded to above. All of these samples have been well

characterized either before or in this study. Also, this study is intended to help elucidate the

effect composition and crystallization history have on the vibrational spectra of the

phlogopites. The muscovite sample, a dioctahedral mica, provides a good comparison for

the trioctahedral phlogopites. Although biotite is also trioctahedral, it has higher iron

content than most of the phlogopites.

Unlike the eloborate sample preparation of thin-sections for volcanic glasses this

study required minimal sample preparation. Grains free (or almost free) from any adhering

phases were picked, cleaned and dried in an oven before analyses. All the samples studied

were already thin plates (~1O-300 urn).

As noted in the Introduction, infrared spectroscopy has been utilized extensively in

the study of micas and other hydrous minerals largely due to its high sensitivity of the

hydroxyl group and its environment. Raman spectroscopy, on the other hand, has been

used sparingly (Tlili et. al., 1989). However, to obtain complete vibrational information, it

is necessary to use both Raman and infrared spectroscopy because these are complementary

techniques. The present work, therefore, is intended to show that micas are indeed good

samples for Raman spectroscopy. The greater sensitivity of infrared spectroscopy is also

used to determine if there are any features in the infrared spectra of micas which reflect

differences under various conditions which they formed.

139Table 13 gives the band positions and assignments for the Raman and infrared

spectra for all the samples studied. Typical Raman and infrared spectra are shown in

Figures 44-50. In order to simplify the discussion the spectral data are sub-divided as

follows:

B. Raman low wavenumber region =130 - 800 cm-1

C. Infrared region =1600 - 2000 cm- 1

D. Raman and infrared hydroxyl stretching region =3000 - 4000 cm- 1

E. Near-infrared region =4000 - 4500 cm- 1

B. Raman low wavenumberregion (130-800 cm-1):

Four bands consistently observed in the Raman spectra of all the phlogopites

studied occurred at approximately 160, 192,353 and 681 cm- 1(Figure 44). The band at

160 cm-1 was observed in all the spectra even though it was weak. This was, however,

not the case for spectra reported by Tlili et. al., (1989) who rarely observed this band. Loh

(1973) assigned this band to the vibrational mode Eg (u2 deformation mode ofM06) for

point group 0h (Herzberg, 1945). On the other hand, a band at -100 cm-1 which was

observed by this same group wasn't seen in any of the spectra in this work. Loh (1973)

noted that this band is frequently hidden by the noise and the laser line. However, this band

wasn't observed in this study because the spectral region was limited to -130-800 cm-1

inorder to avoid the laser (Rayleigh) line from being detected. The diode array detector can

get permanently damaged if the laser line is allowed to hit the diodes for an extended period

of time. Ishii et. al., (1967) had suggested that this band was due to interlayer cations.

However, Loh (1973) and Rosasco and Blaha (1980) showed that it was observed even in

the absence of an interlayer cation. Tlili et. al., (1989) noted that the band can be strong in

both di- and tri-octahedral micas and its position varied with the chemical composition of

micas.

140

s = strong, W=WI.lkM =medium, Sh =shnuldcrDr = broad, VW =vcry weak

Muscovrn~ morrrs CK31 nANDCM·I RAMAN IR RAMAN IR RAMAN IR ASSIGNMENTH)()

ISO lS9(W) IS9(W) lS8(W) M06INTERNAL211O 191(8)250 266(S)

0·11-0 ISOSCELES-300TRIANGLH350 344(5) 360(M) Si-O-Si BENDING400 41S(M) 400(W, Br)

450500550 S37(W,8h)600650 638(W) 667(M,Sh)

Si·Q.Si STRETCIIING700 708(5) 686(M,5h) 684(5) O·H L1DRATIONAL750 757(W)FREQUENCIES800

1150900950

Si-O·AI1000ANTI·SYMMI.nRIC

STRETCII1600 162.S (w)1650 1652(S,Dr) I632(S) COMBINATION17001750

ALUMIN051LlCATH1800 18Ifl(S) 1816(Sh) IlI08(W, Sh) COMOINATION MODES1850190019502()()()

2008(W)2050

203S(w) COMBINATION2150

355036003(,25 3625(S) 3625(S) 3MO{S,nr) 3625(M) 3610{S,l\r) 3600(5)3650

VO-II- LOWER.3675I·REQlll~N<:Y3700 3699(M)

370537103715 37IS(M) 3715(M) 3714(S) O·H STRETCIIING

41004150

VO-II (COMnINATION)4200419S(W) +Si-{)·Si BENDING4250

431K)4291(W) + 0-11 UBRATIONAL4550 .

141

Table 13. (Continued) Band positrons and assignments for the Raman and infrared spectraofmjcns

S - strong, W-wL::lk

M = medium, Sh = shoulderBr = broad, VW =vcry weak

eM·1 CK32 K3 IlD308ll llANORAMAN IR RAMAN lR RAMAN IR ASSIGNMENT

100150 15R(W) I58(W) 158(M) M06IN·mRNI\!.200 193(S) 1,)3(S) 190(M)250

0·11·0 ISOSCELES.300

TRIANGLE350 354(W) 353(W) 353(W) Si·O·Si BENDING400450500550600650 678(S) 682(5) 682(M) Si·O·Si STRETCI lING700 o.n LI13RATIONAL750

I'REQUENCIESROO 744(W)850 786(S)900950

Si-O-AItoooANTI-SYMMETRIC

STRETCII1600 1604(S)1650 1662(5) 1632(W) 1632(M) COMBINATION1700 1712(W) 1710(Sh)1750

ALUMINOSILCATE1800 1803(W) 1R07(W,Sh) COMBINATION1850

MODES19001950 1947(5)2000 2014(W) 2023(W) COMBINATION2050 2072(W. Br)2150 2143(W,llr)

3550 3547(W.llr)3CiOO36253(,50 3610(M) VO-II -1.0WI'.R -31\75 3667(W) FREQUENCY3700 3693(W) 3(,')3(S)370537103715 3715(5) 3714(S) 3714(M) 3714(5) O-H STRETCHlXG

41004150 VO·II (COMBINATION)4200 4199(V WK) 4299(W) 429R(W) + Si-O-Si BENDING42504300 4300(W) 4310(W) 4203(W) + 0·11 LIllRATIONAL4550

-

142

Table 13. (Continued) Band positions and assignments for the Raman and infrared spectraof micas

S =strong, W=wC<lkM =medium, Sh = shoulderBr = broad, VW =\'cr}' weak

CM·I \lD3655 BD3654 \103076 BANDRAMAN IR RAMAN IR RAMAN IR J\SSIGNMENT

100150 I55(W) 158(W) 159(W, Br) M06 INlT:RNAL200 191(5) 191(VW) 191(M)250 0·11-0 ISOSCELES.300 TRIANGl.E350 354(M) 358(V W) 360(M) Si-O·Si \lENDING400450500550600650 6R2(S) 6R2(W) 360(M) Si-O-Si STIUITCllING700 O·H L1URATIONAL750 FREQUENClr:..<:jROO850900950 Si·O·AI1000 At'::rI·SYMMETRIC

STRETCHl(iooIC,50 1631(1\.1) 1632(M) 1631(S, Br) COMBINATION1700 1708(W,Sh) 1705(W) 1705(M)1750 ALlJMINOSILCATE1800 1813(W,Sh) I797(Sh) COMBINATION1l!50 MODES190019502000 2030(W) 2024(W) 2014(W) COMIlINATION20502150

355036003625 3541(W) 3540(W)3650 J605(V W) 3602(W) 3543(W) VO-II- LOWER·3675 FREQlJENCY3700370537103715 3715(S) 3715(S) 3715(S) 3715(S) 3715(M) 3714(S) 0-11 STRETCIIING

41004150 VO-II (COMIlINATION)4200 4295(W) 4197(W) 4202(W) + Si·O-Si BENDING42504300 4207(W) 4298(W) 42<)('(W) + 0-11 LIl3RATIONAL4550

lhble 13.ofmicjls

143

CM-I FRIl493 FR13483 K8 BANDRAMAN IR RAMAN IR RAMAN IR ASSIGNMENT

100150 I59(M) 159(W) 159(W) MOOINTERNAL200 I93(M) I87(W) I91(W)250 O-H-O ISOSCELES-300 TRIANGLE350 355{W) 356(W) 356(W) Si-O-Si BENDING400450500550600650 681(S) 681{M) 682(S) Si-O-Si STRETCHING700 O-H LIBRATIONAL750 FREQUENCIES800850900950 Si·O-AI1000 ANTI-SYMMETRIC

STRETCH16001650 1635(W.Br) 1635(M) 1632(M) COMBINATION1700 1716(W.Dr)1750 J754(W.Sh) ALUMINOSILCATE1800 J797(W.Dr) I797(W,sh) 1807(W,Br) COMBINATION1850 MODES19001950 1933(W)2000 2029(W,Tlr)2050 2056(W.Br) COMBINATION2150

355036003625 3544(VW,Dr) 3541(W)3650 3667(M) VO-II - LOWER-3675 3695(S) FREQUENCY3700370537103715 3712(M) 3714(M) 3708(M) 3714(S) 3714(M) 3716(S) O-H STRETCHING

41004150 VO-H (COMBINATION)4200 4201(W) 4204(W) 4199(W) + Si-O-Si BENDING42504300 4298(W) 4289{W) 4296(W) + O·H LIBRATIONAL4550

S =strong, W-weakM =medium,Sh =shoulderBr = broad, VW =very weak

eM·1 K17 B/\NDR/\M/\N IR /\SSIGNMI:NT

lOO150 15R(W) M06IN11,RN/\L200 IR1(VW)250 0-11·0 ISOSCELES·300 TR[/\NGLE350 349(V W) Si·O·Si BENDING400450500550600650 (,!\O(W) s..o.s: STRETCIIIN(i700 0·11 UBR/\TION/\I.750 I'REQlJENCIliSSOOS50900950 Si·O-/\11000 /\NTI·SYMMETRIC

STR ETC I I1600 1630(M) COMBIN/\T'ON1(,501700 1738(Sh,M)1750 i\I.lJMIj';OSILC/\TI:IKOO COMflIN,\TJON1850 MODESI'JOO195021100 1'J'J2(W) COMBIN/\TION205021511

.l5511

.l60())6251650 VO·II . LOWER.)675 FREQUENCY37003705 3704(5) 3712(M)37103715 0·11 STRETCIIING

41004150 4197(W) Va-II (COMBIN/\T[ON)4200 + Si-O-Si BENDING4250 42S9(W)·1100 + 0-11 LIBR/\TION/\L4550

5 =strong, W-weakM =medium, Sh = shoulderIIr = broad, VW =vcry wcak

144

tra

2000 -I rlcc-.:;

A~

I~en

IIcQ,)~

~ccC

=' sr,

.enc:ec

r",QJ1000 -l I \ Ir•.+J

1"1-IoJ0(,)

U')

.....

.p­Ut

200600Romon Shift

400(cm-1)

Figure 44. Raman spectrum of a phlogopite sample (BD3088) in the 145-800 cm- 1 region.

O I' I 'i i

800

1500

\C\Crl

200

A+~

CfJ 1000I ccc 0

c;j) r--..........C

CJ')

C·C

I II <r.Q) -.......... ~..........0u 500(/)

r-­<r,r-

:1'600 400Raman Shift (cm-l)

Figure 45. Raman spectrum of a muscovite sample in the 145-800 cm- I region.

sr,

.....~Q\

N-r-r<".

I ~;:;:;~

f.4-

Z?

(l)c:J~cd.c r--,

.... ::'\-0 "'!"

l1.I

t.c< =-cc

rlI -.2

-:!:J4000 3000 2000

Wavenu:rnbers (c:rn-l)Figure 46. Infrared spectrum of a phlogopite sample (KI7) in the 1000-4500 cm-1 region.

J Io iii

1.5 T'----------------------.....,

ir:("I

I-.::;re:

J1

~0s:::1=..a~000

..a

..",

.5

I:'\TERFERENCE FRI~GES

\o~ ~(':::">~~ - I

4000 3000 2000WavenuInbers (cID.-1)

Figure 47. Infrared spectrum of a muscovite sample in the 1000-4500 cm- 1 region.....~00

4000 I i

3714

3700WavenuID.bers (cID.-l)

Figure 48. Raman spectrum of a phlogopite sample (8D3088) in the hydroxyl stretchingregion (3600-3800 em-i).

tEl

~ 2000~o

t.>

o --t.\..../'--\.....j.\./\.r'''-v\ .\~\/\~r-J

3600

-~

en~a::sou

4000

2000·

o

3625

3700 3600WavenuInbers (CIn-l)

Figure 49. Raman spectrum of a muscovite sample in the hydroxyl stretching region

(3500-3750 cm- l).

3500

­LIlo

CJCJ~~

..0s..orn

.a<

1

.5

3il91~

3GOI

k"""

4000 3000 2000Wa,·cnurnbers (ern-l)

Figure 50. Infrared spectrum of a biotite sample (RHP-l) in the 1000-4500 cm-! region.-VI-

152

The -192 em-1 line was observed for nearly all the phlogopites and biotite but not

for muscovite. The intensity varied from strong to weak. Tlili et. al., (1989) noted, from

polarization studies, that this line is strong in dioctahedral micas when the electric field is

orientated perpendicular to the cleavege plane, but becomes weak when the electric field

was parallel to the cleavage plane. They also observed the peak position to increase in

wavenumber with increasing Al(IV) and with AI(VI) in both di- and tri-octahedral (K,Na)-

micas.

From polarization Raman studies on the 196 cm- 1 (A1g) peak, Loh (1973) showed

that the tensor component (zz) was much higher than (xx) or (yy) and hence indicated a

distortion occurs in the direction normal to the sheet, since oxygens in M06 are, as

nonbridged 0, bound to their respective Si in this direction. Among the micas measured,

the Raman spectra of clear phlogopite were most consistent with S6 symrne~.

Loh (1973), from his far infrared and Raman studies observed two strong bands

below 180 cm- 1 in the infrared. These bands shifted progressively toward higher

frequencies as the sample varied from biotite, where the bands were at 84 cm-1 and 142

em-I to 92 cm-1 & 156 em -1 and 92 cm-1 & 161 em -1, for brown and clear phlogopite,

respectively and 108 cm- 1 and 165 cm- 1 for muscovite. Raman spectra for these samples

gave three predominant bands below 210 cm-1( at -106,165 and 196 em-I). Unlike the

infrared bands, the Raman peaks were considered to be the same for the brown and clear

phlogopites. Biotite showed no Raman bands in this region. Overall the infrared bands

were observed to be much more sample dependent than the Raman bands. These low

frequency bands were attributed by Loh (1973) to the internal vibrations of the octahedron

M06' with site symmetry 0h' where the M is the octahedral cation, for example Fe2+ in

biotite, Mg2+ in phlogopite and A13+ in muscovite. The surrounding oxygens are: four

from the nonbridged oxygens on four SiO4 tetrahedrons and two oxygens from two OR

ions. The anti-symmetric stretch and bending modes of MO 6 are infrared active (Herzberg,

1531945) and require the displacement of both M and O's and are therefore more sample

dependent because of substitution of Mg2+ by Fe 2+ or A13+ ions. For example, u3 and

u4 in biotite, which has heavy Fe 2+ ions at octahedral sites, are lower than the

corresponding frequencies in muscovite, where the octahedral ion, A13+ is lighter. The

symmetrical stretching Raman active modes on the other hand, require the displacement of

O's only (Herzberg, 1945) and hence vary only slightly from phlogopite to muscovite.

Loh (1973) and Tlili et. al., (1989) observed a general increase in Raman

frequencies with decreasing bond lengths. This was attributed to the greater charge of A13+

compared to Mg2+ reducing the mean octahedral cation-oxygen bond length.

In the frequency region 200-300 cm- l, only muscovite exhibited a strong vibration

(Figure 45). This Raman band at 266 cm-1 has been attributed to internal vibrations of the

isosceles triangle O-H-O by Loh (1973), where O's are the nonbridged oxygens from the

neighboring Si04 tetrahedrons and H is the dangling hydrogen on the OH ion tilted toward

an octahedral cation vacancy. This O-H-O isosceles triangle can only occur in the

dioctahedral mica, where one third of the octahedral cation sites are vacant. The orientation

of OH ions in muscovite is known to be different from that in phlogopite. In the former the

OH is almost parallel while in the latter it is almost normal to the cleavage plane. The

infrared absorption at the OH stretching frequency uO-H =-3625 cm- l in muscovite is,

therefore, strong, while for phlogopite IR band, which appears at uO-H = -3714 cm- 1, is

weak. However, its intensity is increased as its cleavage plane is tilted toward the infrared

beam (Vedder 1964; and Vedder and McDonald 1963).

There are three vibrational modes in the OHO isosceles triangle (Herzberg, 1945),

which are all infrared and Raman active. The other two bands were not observed because

they were probably too weak. Tlili er. al., (1989) observed this band at -270 cm! in all

dioctahedral micas. They observed the position of this band to decrease with increasing

154AI(IV) or increasing Al(VI) in dioctahedral K-micas. Although they also observed this band

in some trioctahedral micas, its dependence upon the Al content was less clear.

The band at -353 cm- 1 varied in position and intensity (weak to medium, Table 13).

This band was observed at -360 cm- 1 for the Raman spectra of (K,Na)-trioctahedral micas

by Tlili et. aI., (1989) and was attributed to the Si-O deformation (SiD4 bending) vibration.

Loh (1973) has assigned the band at -415 cm- 1 for muscovite to the overlapping of the

(OH) libration (i.e., hindered rotation) and Si-O deformation vibration bands (Figure 45).

The strong band at -708 ern-1 in muscovite and at -681 cm-1 in phlogopites is assigned to

Si-O-Si deformation vibrations. Tlili et. aI., (1989) attributed bands at 702 cm-1 and 679

em-I to the deformation vibration of Si-O-Si for muscovite and phlogopite, respectively

whereas, the -640 em -1 and -654 em -1 bands for muscovite and phlogopite, respectively,

were suggested to result from Si-D-AI deformation vibrations (Tlili et. al., 1989). Their

intensity was observed to increase with increasing AI(tot) (AI(IV) + Al (VI» in trioctahedral

micas. The other bands observed for muscovite at 638 and 757, cm- 1 are attributed to Si-O

(Si04) and AI-O-Si stretching modes, respectively (Figure 45).

C. Infrared region (1600-2000 cm- 1)

The multiple bands (1632,1738 and 1992 cm-I) observed in the 1600-2000 cm- 1

region of the infrared spectum of K17 (Figure 46) may be attributed to alumino silicate

combination modes (George Rossman, private communication). Vedder (1964) observed

these bands, but since they were very weak he didn't give any assignment for them. Since

the assignments are unclear they will not be discussed any further. The band observed at

-1635 cm! for phlogopites and biotite appeared at the same position as the bending mode

of molecular water (Table 13). To determine if this band was due to the presence of

molecular water incorporated within mica, a high temperature study was undertaken. (see

Appendix F)

155

D. Raman and infrared hydroxyl stretching region (3000-4000 em-I)

Vedder (1964) classified the hydroxyl stretching bands into three main types: N­

bands (normal), due to OH groups bonded to three octahedrally coordinated divalent

cations (OH bonded to 3Mg); I-bands (impurity), due to OH groups bonded to two divalent

and one trivalent cation (OH bonded to 2Mg, Al or Fe); and V-type (vacancy), due to OH

groups adjacent to an octahedral vacancy. Hydroxyl-stretching wavenumbers were

observed to fluctuate as a function of the mica composition.

The infrared spectra, in the hydroxyl stretching region, are all very similar for the

phlogopites. A single band was observed at -3714 cm- l(N-type) for all these samples.

This has been interpreted as an N-type stretching vibration of OH influenced by three

divalent cations by Vedder (1964). Robert and Kodama (1988) observed the shift of this

band as a function of the total aluminium content. All the phlogopites studied here have

very similar aluminium content (8.61-12.75 wt, %) which is reflected by the hydroxyl

band position being constant for all the samples studied. The Raman and infrared spectra of

all phlogopites and muscovite (Figures 48 and 49) both gave a single hydroxyl band at the

same frequency. This has been previously observed by Robert and Kodama (1988). The

intensity of the hydroxyl band in the infrared and Raman spectra for muscovite (Figures 47

and 49) was stronger then the same band observed for phlogopites (Figures 46 and 48).

The hydroxyl band in the infrared spectra of all the phlogopites is shown to be strong in

Table 13, because it is strong relative to the other bands observed in the same spectrum..

The muscovite band was not only strong but it was also shifted to lower frequency,

-3625cm- 1 (Figure 47). Vedder (1964) assigned this band as a V-type stretching vibration

of OH close to two trivalent cations and a vacant site. This is attributed to the difference in

orientation of the hydroxyl in muscovite relative to phlogopite. The biotite sample gave a

broad, complex envelope (I-type). High iron content may lead to a variety of hydroxyl

sites, rather than a single site observed for the other two types of micas. Another biotite

sample (RHP-l), obtained from the University of Colorado gave a better defined spectrum

(sharp doublet; Figure 50) attributed to two joined polytypes.

156Some infrared spectra also showed other bands in this region, varying from

generally weak to medium in intensity. These bands are attributed to other phases present

within the grain. (Matson et. al., 1986) These extra bands may be due to hydroxyl

stretching from other hydrous phases present in the mica grain (i.e., chlorite). However,

these bands didn't interfer with the Raman spectra where the laser beam diameter was much

smaller at approximately 2-3 11m. The infrared utilized a -600llm aperture, hence there is a

greater likelihood of interference from other phases.

E. Near-infrared region (4000-4500 cm-1)

The bands observed at approximatly 4200 cm- l and 4300 em-I have been attributed

to combination modes between uO-H Stretch (-3714cm-1) + (496 and 595). These two

latter bands are due to Si-O-Si bending and uO-H libration (i.e., hindered rotation) modes,

respectively. Their intensities are weak as normally expected for combination modes.

Similar bands have been previously observed by Vedder (1964) on his work on the

correlation of infrared spectra with chemical compositions of micas.

157APPENDIXE

The effect of hi~h pressure on the infrared spectra of micas

A complete knowledge of the chemical composition and phases of materials present

at high pressure is fundamental to our understanding of the interiors of planets. Hydrogen

bridges, O-H---O, are important structural elements in certain silicate phases. The energies

of OH-vibrations are determined by the strength and geometry of the hydrogen bridge.

Changes in the structural environment can be detected through changes in the energies of

the OH-vibrations when measured by infrared spectroscopy (Rossman, 1988). Infrared

spectroscopy, due to its high sensitivity for hydroxyl groups, has played a major role in the

investigation of micas at ambient conditions. Valuable crystallo-chemical information can

also be obtained from a knowledge of the P-T-X dependence of OH-bond energies.

The influence of temperature on the stretching vibrations, uO-H' of various OH­

bearing silicate minerals has been studied by Freund (1974) and Aines and Rossman

(1985). The influence of composition of silicates on the uO-H bond has been studied by

numerous people (e.g., Rossman, 1988). Although Langer et, aI., (1979) studied the

influence of pressure on hydrogen bonds in opals, there is limited infrared data on the

systematic change of the uO-H bond with respect to increasing pressure. Winkler et. al.,

(1989) studied the affect of high pressure upon the hydroxyl stretching bands of zoisite and

Kruyer et. aI., (1989) have studied its affect on brucite (Mg(OHh) and portlandite

(Ca(OH)2)' Both found that the hydroxyl stretching band shifts to lower energies as the

pressure is increased. This has been interpreted to result from an increase in hydrogen

bonding that can occur between the proton of the hydroxyl and a close electronegative atom

(oxygen). As the hydrogen bond strength increases, there is a concomitant weakening of

the O-H bond, resulting in the uO-H band shifting to lower energies.

In this work, the effect of pressure upon the hydroxyl stretching band for a

phlogopite (K8) and muscovite sample has been investigated using Fourier transform

infrared spectroscopy (FTIR). The pressure was applied using a diamond anvil cell (DAC),

which is capable of generating pressures in the megabar range. During the past several

158years advances in both vibrational spectroscopy (Raman and infrared), as well as advances

in high pressure DAC designs have made it feasible to investigate the structure and

vibrational properties of materials (especially minerals) at very high pressure. The design

and usage of DACs have been reviewed in several recent publications (Ferraro, 1986;

Sharma, 1989a,b).

1. Experimental

Figure 51 shows the basic configuration of the diamond anvil cell. A sample in a

punctured metal gasket mounted between the flat parallel faces of two opposing anvils, is

subjected to pressure when the anvils are brought closer together. Most of the DACs can

be used for generating pressure s; 350kbar. The cell used for this study is based on a

design of Mao and Bell (1978a,b) (Figure 52) for generating pressures exceeding 1

megabar (Mbar). The sample was loaded under a binocular microscope by first seating the

600Jlm molybdenum gasket on one of the diamond flats with the hole in the centre. The

sample was then placed into the gasket and the rest of the space was filled with KBr

powder to provide hydrostatic pressure. A few grains of ruby powder were also embedded

for measuring pressure. The ruby grains provide a very elegant way of measuring pressure

in-situ by measuring the red shift of the R1 fluorescence line of the ruby crystal (Piermarini

et al., 1975). The shift is almost linear (Figure 53) up to 300kbar at a rate of 0.365 A /

kbar or 0.753 cm! kbar-1. The following equation proposed by Mao et. al. (1978) can be

used for estimating pressure from the R1 ruby line shift over an extended range:

P (kbar) = 3808 [ CA/AO)5 - 1] (35)

where A o(nm) is the wavelength of R1 line at 1 bar and A(nm) is the wavelength at

pressure P. The diamond anvil cell was mounted upon a special stage within the sample

compartment of the FTIR. A set of two KBr lens were used, one to focus the incoming IR

beam and the other to collimate the exiting radiation upon the detector ( Adams and Sharma,

(1977); Figure 54). For each pressure increment and fluorescence measurement the cell

had to be removed from the FTIR compartment. So, each time the cell was reinserted into

the FTIR companment , although the signal was optimised before spectral accumulation, a

1 F'orce 1

159

f t t 1

Figure 51. Basic configuration of the diamond anvil cell.

160

Z .,conlum \hll'"'

Tu"Q~'e" carbide,",olf =vhnder

5 L~",!'

6 Ma ..·' body

1 0', ",e ~C:'~.

e B~ll~V.II!

.o'''~r~

o!

TUf"lq,ten cartJ,de"'olf cyhttde'

I C,lt"der

2 P·s'on3 !''''u,t tlloC"4 ~uIC'u""

50""DI,Qcsk.'

B

A

Hordef"led steel :>'5ton, locoed. flt'ed ,ntohorde",d slee' cy""der

Figure 52. Schematic diagram of (A) Mao-Bell (M-B) diamond anvil cell and (8) M-B cellpiston-cylinder details.(after Mao and Bell, 1978a,b)

161

;:;,0

~ \20....a::::::lV1V1....a::0-

80

RUBY R. WAVELENGTH SHIFT ..\A IAl

Figure 53. Pressure calibartion of ruby R1 fluorescence Iine to -160 Kbar (after Block andPiermarini, 1976).

Figure 54. Ray diagram of refracting beam condenser and DAC for infrared measurements(after Adams and Sharma, 1977).

....~

163slight change in positioning of the DAC will result in different sized sample areas being

subjected to the IR beam.

All spectra were collected at room temperature with a Perkin Elmer 1720X FfIR

spectrometer using a globar source and DTGS detector. Spectra were obtained in the 400-

4000 cm- 1 region at a resolution of 4 em-I. A mirror speed of 0.5 em/sec and 500 scans

were sufficient for muscovite. However, the unfavourable O-H orientation for phlogopite

relative to muscovite meant that the hydroxyl stretching band intensity was much weaker.

So, the number of scans necessary to obtain a reasonable spectrum had to be doubled to

1000 scans for phlogopite. Due to slight differences in pressure within the gasket, the

pressure reported here is an average of a couple of points (± 10 %).

2. Results

Muscovite spectra were obtained from 1 bar (1 atm.) up to -200 kbar (200,000

atms). Table 14 gives the pressures studied plus the position at full width at half maximum

(FWHM) for each pressure measurement. Representative high-pressure spectra recorded at

pressures up to 201 kbar are shown in Figure 55. The band at 1 bar is positioned at 3630

em-I. A shift to lower energies is observed with increasing pressure. Also, as seen in Table

14 there is a simultaneous increase in FWHM as pressure is increased.

The phlogopite data are sumarized in Table 15. Representation high pressure spectra

for phlogopite K8 are shown in Figure 56. The data show a general increase in FWHM.

Although, the shift in the hydroxyl band position isn't as dramatic as that observed for

muscovite, the shift can be semi-qualitatively, as a consequence of peak broadening,

thought to increase with increasing pressure. Previous studies by Winkler et. al., (1989)

experienced a similar problem where they determine the peak position by halfing the

distance between the flanks. This method was applied here by using a spectral analysis

program (Spectra Cal, Search Arithmetic S2.12, Galactic Industries Corp., 1988). The

background which is necessary to measure the FWHM is in most spectra somewhat

ambiguous. These FWHM values were measured manually from a hardcopy of the

spectrum, with a pencil and ruler. Because these experiments were undertaken only once

Table 14. The effect of pressure upon the position and FWHM of the hvdroxyl band for muscovite

Pressure (kbar) Position (em-I) FWHM (em-I)

1 3630 6110 3622 5327.3 3622 5755.3 3609 6671.3 3612 5792 3612 74110 3616 74134 3605 82168 3605 86

184 3606 90

201 3590 94

......~

Figure 55. Infrared spectra to show the effect of pressure upon the positions of thehydroxyl stretching band for muscovite.

C\)g

~..aJ.toQJJ

~

3800

169 KBAR

110 KBAR

3800If'avenulDbers (cUl-l)

3400

-0\VI

'Thble 15. The effect of pressure on the position and FWHM of the OH band for phlogopite (K8)

Pressure (kbar) Position (em-I) FWHM (em-I)

1 3715 37

3.3 3717 3914 3719 50.941 3722 50.756 3727 55.3109 3728 73.7145 3729 83.8

-0\0\

C)o~CIS

..0Moen

..0~

145 KBAR

] BAR

·1000 3800Wavcnuxnbcrs (em-I)

Figure 56.Infrared spectra to show the effectof pressure on the position of the hydroxylstretching band for a phlogopite sample (K8).

3600

.-~

168the magnitude of the errors in the band position and FWHM cannot be stated. However,

the variability in the pressure is approximately 10% for each pressure measurement.

3. Discussion

The two types of micas studied here, dioctahedral (muscovite) and trioctahedral

(phlogopite, K8), have very different orientations of the O-H bond. These hydroxyls are

part of the plane of oxygens which is shared by the tetrahedral and octahedral sheets. As a

result, the hydroxyls lie between two sets of cations, the octahedral (Y) sites on one side

and the tetrahedral (Z) and interlayer (X) sites on the other. Since the Y sites are closest to

the hydroxyl, the electrostatic repulsion will favor an OH orientation which maximizes the

Y-H distances. In either a trioctahedral or dioctahedral structure, the maximum distances

will result in an O-H directed away from the octahedral sheet toward the large ditrigonal

cavity of the tetrahedral sheet. Opposing this orientation is the repulsion from the interlayer

and tetrahedral cations which are further away but which are more numerous and

commonly have larger charges. Thus, the OH orientations represent a balance between the

repulsions from these two sets of sites. This is realised by the position and also the band

intensities of the respective bands. The hydroxyl group is almost perpendicular to the

cleavage plane in the phlogopite sample, whereas it is tilted from the normal to the cleavge

plane for muscovite. The muscovite OH band appears at -3630 cm- 1 at 1 bar and the

phlogopite band appears at -3715 em-I. It is clear that the OH in muscovite is experiencing

more hydrogen bonding, due to its lower band position. The relative intensities of the two

bands are very different. In order for the infrared radiation to interact strongly with the

hydroxyl bond it must be parallel to the hydroxyl bond axis. For phlogopite the incoming

radiation is almost perpendicular to the O-H bond and consequently the interaction is weak

resulting in a low intensity OH band.

When muscovite is subjected to higher pressure the hydroxyl band position

decreases (Figure 57). The hydroxyl group is deformed such that it becomes more coplanar

with the cleavage plane. The dangling hydrogen gets even closer to the neighboring

oxygens resulting in an increase in hydrogen bonding. The corresponding O-H bond

weakens and hence the band position is shifted to lower energy. On the other hand the D-H

300200100

3640 I I

3580 I ' I iii Io

---.~

3620~.--Z,....

3610-t:en-.-'e,

36000Z-e:::

3590

Pf~ESSl'RE (KJl.-\R)

Figure 57. Relationship between pressure and hydroxyl band position for a muscovitesample.

.....$

170bond in phlogopite is almost perpendicular to the cleavage plane. When you subject

phlogopite mica to pressure, the O-H bond is shortened, mainly as a result of the interlayer

cations which strongly repel the hydrogen on the hydroxyl group. As the pressure is

increased, the repulsion between the hydrogen and interlayer cation increases, the O-H

band position, therefore, shifts to higher energy (Figure 58). This study, instead of using

orientation (Serratosa and Bradley, 1948 and Bassett, 1960), has been able to determine the

orientation of the hydroxyl group in phlogopites and muscovites by using pressure.

Unlike, the muscovite study where the band position changed almost linearly with

pressure, for phlogopite the position of the OH band appears to change non-linearly in the

pressure range 1-80 kbars thereafter the band position stays almost constant to 150 kbar

(Figure 58). This observation may be attributed to the difference in compressibility of the

tetrahedral and octahedral layers. The pressure dependence of the OH frequency is an

important parameter to consider for phlogopites where the hydroxyl group is almost

perpendicular to the cleavage plane. The application of pressure would have a direct effect

on the interaction of the interlayer cation and the hydrogen of the hydroxyl group.

The FWHM for both micas increases significantly with increase in pressure (Tables

14 and 15). Figures 59 and 60 show the effect of pressure upon the hydroxyl bands of

muscovite and phlogopite, respectively. In both instances the FWHM increased almost

linearly with pressure. Previous work has attributed the widening of vibrational bands to an

increase in anharmonicity (Hertzberg, 1945). It has been stated that if the FWHM

increases there must be an increase in anharmonicity, Moon and Drickamer (1974)

attributed the ubiquitous increase in FWHM of the O-H vibration to steepening of the

configurational coordinate potential well. This also indicates increased bond anharmonicity,

in that the vibrational frequency and therefore the quadratic contribution to the effective 0­

H potential is decreasing with increasing pressure. Consequently, significant anharmonic

contributions are required inorder to steepen the potential. Alternatively, the peak width

may reflect a broad distribution of O-H distances resulting from the increase in disorder, of

the hydroxyl groups, as the pressure is increased.

With the help of the results of Nakamoto et. al., (1955), who empirically derived a

relationship between the energy of the stretching vibration and the 0---0 distance in

3730 ' i

.­-I

~U--zo-~ 3720-rJ1o~

QZ<~

2001003710 -I iii I

o

PRESSURE (KBAR)

Figure 58. Relationship between pressure and hydroxyl band position for a phlogopite(K8) sample.

-.I

100

--I

~o

90-~;;.J

~-><:-e 80~r;-....J<-.....~

70

--Eo-Q-~ 60....J....J;;.J~

50~0 100 200 300

PRESSURE (KBAR)

Figure 59. Relationship between FWHM and applied pressure for the hydroxyl stretchingband on a muscovite sample.

.....-...ltv

200100

PRESSURE (KBAR)

90 i :> i

80

70

50

60

30 -I iii Io

,-.,~

I

:;U---:;;::)

~-x~

:;LI:......:l~

:c~:cEo­~-~....:l....:l;::)LI:..

Figure 60. Relationship between FWHM and applied pressure for the hydroxyl band for aphlogopite (K8) sample.

-..JW

174hydrogen bonds, Winkler et. al., (1989) were able to calculate a 0---0 distance of 2.76 A

at lbar and an 0---0 distance of 2.68 A at l00kbar for zoisite. This distance at lbar was

found to be in agreement with crystallograhic data. It was able to interpolate from

Nakamoto et. aI., (1955) results, for muscovite, an 0---0 value of -3.07 A at 1 bar

and an 0---0 distance of -3.03 A at 200 kbar. It was. however, not possible to use their

data for phlogopite because the highest frequency they recorded was 3700 em-I.

The DAC had to be taken out of the spectrometer for each pressure increment and

fluoresence measurement. Therefore, exact alignment changed with each experiment even

though great effort was made in trying to optimize the energy throughput for each pressure

studied. Intensity comparisons between spectra are only semi-quantitative due to the

pressure induced change of the refractive indices of the sample, in contrast to the nearly

constant refractive index of the diamonds. The siginificant amount of scatter in the band

position with pressure is a consequence of the inhomogeneity in the pressure experienced

by the sample (+ 10 %) and the inability to measure the peak position accurately. especially

when the band broadens.

It is interesting to note that these samples were able to withstand much higher

pressures than those at which they were formed. For example, the phlogopites probably

formed at high pressure (20-30 kbar) and the muscovite was probably formed at low

pressure « 1 or 2 kbar) (M. O. Garcia, Private communication). Infact, the hydroxyl band

positions were observed to return to their original values before the application of high

pressure.

175APPENDIXF

The effect of high temperature on the IR spectra of micas

Studies at ambient temperature provide limited insight into the physical properties of

minerals at the elevated temperatures typical to conditions of geological interest. So, there

is a need to undertake high temperature studies inorder to recreate conditions minerals

normally exist under.

As a consequence of the infrared study of micas at ambient conditions (see Chapter

VI) a broad multiple band was observed at -1635 em-1. A band at -1635 cm-1 is normally

attributed to the deformation frequency for molecular water (Herzberg, 1945) and so clearly

signifying the presence of molecular water (in contrast to OH). The presence of this band

indicates the possibility of water residing within micro-cracks or as an alteration clay

product, which contains water molecules. Previously, this band has generally not been

observed. The one publication where this band was reported, the author noted the band to

be very weak, however, neglected to mention its origin (Vedder, 1964). In order to

elucidate the origin of this band a high temperature study was carried out at 1 atm. If this

band is due to the presence of molecular water adsorbed upon the surface of the sample or

as a clay product it would be expected to disappear when the mica is heated to a couple of

hundred degrees.

1. Experimental

This study provided a direct method of studying the effect of temperature upon the

infrared spectrum of the mica, in the 400-4000 cm! region. A copper metal disc, with a

- 600 urn aperture, was specially made such that it was the same diameter as the small

ceramic heater. A coiled nichrome wire was inserted into the interior of the ceramic disc

heater to be used as a heating element. Both the disc and heater, once aligned, were held by

clips within a slot, having the same diameter as the heater, on a specially made metal plate

which inserted into the slot of the sample accessory in the sample compartment of the FTIR

spectrometer. The tip of a Pt-lO%Rh Pt thermocouple was positioned as close to the sample

aperture as possible, to allow for better temperature measurement at the sample. Three

176samples' were chosen for this study: a muscovite, phlogopite (BD3088) and biotite (RHP-

1). The biotite sample was received from the University of Colorado. These samples were

cleaned prior to their analyses. Before sample analyses, a background spectrum was

obtained in the 400-4000 em -1 region. Thereafter, the mica grain was carefully mounted

upon the aperture, with the aid of a high temperature adhesive. Each sample was heated by

-500C increments and the infrared spectrum was obtained, once the temperature stablized.

The temperature was carefully controlled with the aid of a variac and was observed to

fluctuate, ::;± 50C at each spectral measurement. By using a DTGS (deuetrated triglycine

sulphate) detector the effect of blackbody radiation, which is a continuous phenomenon,

would not interfere with spectral accumulation at high temperature measurements because

this detector responds to alternating signals only.

2. Results and discussion

Figures 61-63 show the spectra obtained as a function of temperature for each

sample. There was a decrease in intensity and broadening of the peaks as the temperature

was increased for all three samples. However, the bands for phologopite and biotite

samples returned back to their original positions and shapes when cooled back to room

temperature. The fact that these bands didn't disappear for the phlogopite and biotite

samples imples that they are probably not due to the presence of molecular water. This

conclusion is based on the assumption that the crystal structure of mica has no site for

molecular water and the presence of any adsorbed water would have been removed by

heating to 400 0C (Matson et. al., 1986). Also, we observe a single high temperature H20

release peak during heating in the high temperature mass spectrometer. These bands may

be attributed to combination modes of the alumino-silicate network rather than molecular

water. However, the band for muscovite slowly decreased and finally disappeared at

-3410C. The sample was left over night, after the experiment, to check if the band would

reappear, however, it didn't. Since the position of the band was at ~1700 cm-1, which is

high for molecular water, it may have been a non-hydrated impurity present on the surface.

Cl,)t)

fJ.cSotoIn

~

. 5

o

BACK AT R.T.

200

1800 1600 1400Wavenum.bers (cm.-l)

Figure 61. The effectof temperature upon the infrared spectrum of a phlogopite sample(BD3088) in the~1400-1900 cm-1 region. --J

-J

1 ROO~I TE~IP

2 200 DEG C

3 400 DEG C

4 BACK AT RT

o

1

C)C)

d I~ .5 ~

..ar...oen..a<:

2000 1500Wavenurnbers (crn-l)

Figure 62. Theeffectof temperature upon the infrared spectrum of a biotite sample(RHP-l) in the-1400-2200 cm-1 region. ......

-...l00

DEG. C

138 DEG. C

"---.-------_/200 DEG. C

-----c:,)c;,)

=:~ 1~

oen.c<:

341 DEG. C

BACK AT ROO~I TE:\!P.

o1800 1600 1400

Wavenurnbers (ern-I)Figure 63. The effectof temperature upon the infrared spectrum of a muscovite sample inthe -1400-2000 cm-1 region. ....

-J\0

180Part of the reason why other workers have not reported molecular water bands may

have been because many of those studies used KBr pellets. Since the concentration of mica

sample would be low this band would probably not be visible. Also, the great affinity of

KBr for water would result in a band at -1635 cm! and so mislead the researcher. George

Rossman (California Institute of Technology, personal communication) felt the major

reason why this band hasn't been mentioned before is because it is not of primary

importance. He suggested that the band is either due to combination alumino-silicate bands

or to the presence of chlorite (deformation band attributed to NH4+) which is a mineral

phase commonly found in mica sheets.

This study also provided the opportunity to observe the effect of temperature upon

the hydroxyl stretching band for muscovite (Figure 47). The other two (phlogopite

(BD3088) and biotite (RHP-l» samples had to be thick to observe the broad band at -1635

em-I and consequently the hydroxyl band was totally absorbing. This band was observed

to undergo little shift in position even when using the highest resolution of 2 cm-Iso the

shift is most probably less than 2 ern-1. However there was a slight decrease in intensity.

Previously, Aines and Rossman (1984) from their study on the effect of temperature upon

the OH-stretching region of muscovite reported a decrease in intensity as the temperature

reached -5900C and a shift to lower energy. However, this shift was small until the

temperature reached -590°C. This shift to lower energy was attributed to the lengthening

of the O-H bond as the temperature was increased. There was also a drastic drop in

intensity between 590°C and 757 0C which Rossman (1984) attributed to the dehydration

of the muscovite. Aines and Rossman (1985) noted that the speciation and properties of

minerals containing trace hydroxyl and water changeddramatically between observing them

at room temperature and at temperatures of geological interest. For instance, they observed

no changes in speciation prior to dehydration at 750°C for muscovite whereas, topaz

hydroxyl sites interconvert at 500°e. They also noted that at high temperatures the

181following changes were common: broadening, a shift to lower wavenumbers, and a slight

decrease in integral intensity.

Freund (1974) has discussed the reason for the decrease in intensity of the infrared

band with increasing temperature. Absorption of infrared radiation takes place when the

dipole moment of a given atomic arrangement changes during a vibrational transition. If Jlo

is the dipole moment before the transition, the dipole moment Jl after the transition is given

by:

where 0<4c is the atomic displacement expressed in the normal coordinates Qk. For small

displacements, the intensity of the infrared band is proportional to the dipole moment

derivative (OJ.1l0~). Due to anharmonicity the mean distances between atoms in a solid

generally increase with increasing temperature, making the absolute values of the dipole

moment Jlo larger, but the relative changes (OW0<4c)Qk smaller. Hence, the integral

intensity of the infrared bands of solids generally decreases as the temperature increases.

At the same time, the anharmonicity also makes the bands broader as the temperature

increases.

182APPENDIXG

The effect of low temperature on the Raman spectrum of muscovite in the

hydroxyl stretching region

Most spectroscopic studies of minerals have been undertaken at room temperature,

due to its experimental convenience. Low temperature studies have helped to resolve broad

bands into their individual components, when the broad band is due to multiple bands that

happen to appear close to each other. Kats (1962) studied the spectrum of quartz in the

hydroxyl region. He observed a nearly continuous blur at room temperature tum into a set

of sharp bands at -195 0 C. These bands were attributed to the association of OR with

specific alkali ions such as Na+, Li+, and K+. Wang et. al., (1988) used low temperature

(-190 °C) to reduce the linewidths of the OH bands, for cummingtonite, in an attempt to

resolve the predicted fine structure for the amphibole.

This study will investigate the effect of low temperature upon the Raman spectrum

of muscovite in the hope of elucidating the nature of the broad single band observed in the

hydroxyl stretching region for the Raman spectrum of muscovite. The Raman and infrared

spectra of muscovite are similar in the hydroxyl stretching region (see chapter 6 (dj), Both

give a single strong band at approximately 3625 em-I. This band appears to be broad

relative to the single hydroxyl band observed for phlogopite at -3715 em-I.

1. Experimental

The effect of low temperature upon the Raman spectrum of muscovite was explored

using a commercially available helium cooled cryostate (C'I'I Cryogenics). This apparatus

was used to cool the muscovite sample down to -33 OK with O.loK accuracy. The sample,

a small sheet of muscovite, was afixed upon the cold finger with the aid of adhesive. The

cold finger was held under a vacuum (-5-10 x 10-3 torr) by the aid of a fore-pump to avoid

frosting of the optical window. A sapphire optical window, on top of the cold finger

holding the sample, allowed the laser beam to be focussed upon the sample. The sample

was initially cooled down to the lowest temperature by using a microprocessor-based

183temperature controller connected to a gold-bismuth thermocouple to sense the temperature

at the cold finger. The temperaturecontrolleradjusts the heating load to maintain the desired

temperature at the cold finger. Higher temperatures were obtained by slowly heating the

cold finger and setting the controller to the desired temperature value.

A micro-Raman set up using 1350 geometry was used (see Experimental Methods

section) The Ar+ 457.9 nm laser (Spectra Physics 2020) line was utilized to excite the

sample with an average power of -20mW at the sample. The Raman scattered signal was

collected by an 20X objective and detected using a CCD detector, cooled to a temperature

of -1200C using liquid N2 as the coolant. The hydroxyl stretching region (3000-4000

em-I) was calibrated as outlined in the Experimental Section. The 457.9 nm line was used

for all spectral analyses. This calibration was accurate to within + 1.4 cm-1 when the

Raman signal was dispersed with a 1200 grooves/mm grating. The sample was focused

underneath the microscope and then the laser beam was allowed to hit the sample. A CCD

camera allowed me to observe the sampleon a TV monitor before spectral collection.

2. Results and discussion

The Raman spectrum was limited to the region, 3000-4000 cm-1 (Raman shift).

Only one band was observed at 3625 cm-l at room temperature (Figure 49). This band was

observed to shift to 3631cm- 1 for the lowest temperature of 33°K. All spectra in between

slowly decreased in position, such that the final position recorded at room temperature was

back at its original position of 3625 cm-1(Table 16). The full width at half maximum

(FWHM) did not change during the course of the experiment. The shift in position, which

could be measured with a good deal of certainty due to its distinct sharp maximum, of the

hydroxyl band with temperature was found to be linear (Figure 64). The expected

deconvolution of the broad hydroxyl band wasn't observed, so, it wasn't possible to say,

unambiguously, if this broadness is attributed to the, almost, coalescence of multiple OH

peaks.

Table 16. Effect of low temperature on the OH position of muscovite

184

Temperature (oK) Position (cm- 1)

33.52 3628.5

66 3627.5

100 3627.5

142.5 3627

Room temp 3625.4

300200100

3629 I I

3625 I iii I i Io

.-

.-t 3628I

~U'-'

;Z0""""E- 3627

""""rJJ.0~

Q;Z< 3626=:l

TEMPERATURE (K)Figure 64. Relationship between position and temperature of the hydroxyl stretching bandfor muscovite.

t-ooVI

186The probable cause for this broadening could be attributed to the slightly different

hydroxyl orientations within the muscovite. This would result in slightly different O-H

bond lengths within the structure since there would exist differing amounts of hydrogen

bonding. So, one would observe this distribution of O-H band lengths as a broad band.

The different orientations would also effect the intensities as mentioned in the Introduction.

The almost symmetric band suggests that the majority of the hydroxyl bonds are inclined at

a single angle, with the rest of the angles equally distributed slightly above and below this

value. Unlike ideal structures for muscovite, natural muscovites experience some

subsitution in the octahedral layer which consequently results in slight changes in the

hydroxyl orientation.

187APPENDIXH

I. Hi~h temperaturemass spectrQmetry rnTMS)

Water, structurally bound as hydroxyl ions, is an important component of hydroxyl

minerals such as mica. SQ, it is important to report its concentration as part of the

compositional analysis when possible. Hydroxyl (or water) contents of minerals cannot

be determined using electron microprobe techniques. Water, when reported, is generally

determained as H20+ by heating the mineral to release its water content which is

catalytically reacted to form H2 and measured manometrically. High temperature mass

spectrometry (HTMS) , apart from providing volatile abundances for all volatiles contained

in the mica, has the potential to provide information helpful tQ constrain the site occupancy

of hydroxyl grQUPS within the mineral's structure.

Previously, Matson (1984) and Matson et. al., (1986) used a combination of

HTMS and electron microprobe tQ determine the volatile and major elemental composition,

respectively, of a collection Qf phlogopite micas from south African kimberlites. The

volatile release behavior of all the phlogopite micas analyzed were similar regardless of

their mode of occurrence. Most of the major volatile species in these micas were released in

the relatively narrow temperature range from 1OOQ-11500 C (Figure 65). The water release

peak typically exhibited a IQW temperature shoulder extending below l000oC. In addition

to the major volatile release at -llOOQC, some of these micas produced small-to-moderate

intensity IQW temperature (650-750 QC) volatile release peaks which generally involved

water and/or small amounts of carbon- and sulfur-bearing species (e.g., C02' CO, CH4,

H2S) and tended to be more prominent in (but not limited to) the release patterns of

samples from peridotite nodules. Matson et. al., (1986) have suggested that the IQW release

could result from the presence of additional mineral phases which were not effectively

seperated during the sample picking procedure. Due tQ the presence of alteration products

(chiefly chlorite) which has been observed to dehydrate between 600 and 700 0 C by

differential thermal analysis (DTA), they decided to discard the IQW temperature release

188

(a)

1.6

2.4

3.2

f- 0.8--J

0>>- af-

V>ZWf-

Z (b)0.6

Z0

1300700 900 1100TEMPE RATURE (O()

OL--_--L__....L-__1-_--L__-1-__..I--_--...I.__~

500

0.2

0.4

Figure 65. Mass pyrograms of a phlogopite mica megacryst (FRB483) showing typicalrelease behavior of: (a) H20, F, Cl and (b) CH4 (as CH3+), CO, CO2, (after Matson,

1984).

189when calculating total weight % of each species contained in the mineral. They calculated

the chemical formulas of the micas on the basis of 24-anions according to the method of

Deer et al., (1982).

Despite the water and halogen losses from the hydroxyl sites in the crystal

structures of most of the micas at -11oooc, the identities of individual phlogopite mica

grains were preserved to at least -13000e. After loss of their volatiles the micas appeared

greyish in colour and were quite brittle. In addition, grain surfaces appeared blistered.

(i). Experimental

\blatile contents of the hydrous minerals analyzed in the study were obtained using

the high temperature quadrupole mass spectrometer in our laboratory. Details about the

procedure and instrumental aspects can be found in the Experimental Methods section of

Part 1 of this dissertation. The mica grains were typically larger than lrnm in diameter and

thickness varied from very thin flakes to 0.5 mm or more. However due to the higher

volatile content of muscovite -22 mg was adequate for this analysis. This sample size is

similar to that used for phlogopites by Matson (1984). The muscovite sheet was transparent

unlike the phlogopites studied previously which ranged from light brown to dark brown.

(ii). Results and discussion

Unlike, phlogopites, where most of the major volatile species were released in the

relatively narrow ternperture range, 10OO-11500C, most of the volatiles from the muscovite

were released in the broader interval800-11000C (Figure 66 a,b). The small to moderate

low temperature (650-7500C) volatile release observed for phlogopites wasn't observed in

this study. This is a good reason to assume that the sample was not contaminated with

other phases as was the case for the phlogopites. Table 17 gives the volatile content of the

muscovite sample. Unlike the phlogopites which contained between 3.53-4.86 wt. % total

water, this sample had -2.53 wt, % H20, showed considerably higher abundances for

e02 and eo and no detectable Cl (see Table 17). After heating, the sheet had become

brittle and greyish in colour. It was no longer transparent and the surface appeared

190

(a)

I---I

0>~l-

V)

Z (b)wI-

Z

Z0

200 400 600 BOO 1000

TEMPER AT URE (O()

Figure 66. Mass pyrograms of a muscovite sample showing the release behavior of:(a) H20, HF, F and (b) CH4 (as CH3+), CO, S02' CO2.

Table 17, Elemental analyses for a muscovite and biotite samples (wt.%)

Sample Muscovite Biotite

Si02 45.83 38.93

Ti02 0.16 2,02

Al203 33.65 15.67FeO* 4.52 9.64MnO 0,13 o. 1M!:,7() 1.52 21.31Na20 n.65 O.3RK20 9.27 8.84F 0.52 n,d.CI 0 n.d.H2 O 2.53 n.d.

CO2 0.918 n,d,

CO 0.253 n.d,CH4 0.257 n.d.S 0.0524 n.d.Tot31 100.3

Ionic Formulas Calculated on the Basisof 24 (0, OH, F, CI)

Si 6.315Alz 1,685Aly 3,781Ti 0,017Fe 0.521Mn 0,015Mg 0.312Na 0.174K 1.629F 0.227CI 0OH 1,598

Cation Totals

Z 8y 4.645X 1.803Mg# 37.9

n.d. = not determined

191

192blistered, which Matson (1984) attributed to the high internal vapor pressure of the released

volatile species, and the resulting delamination of the mica sheets.

2. Electron microprobe analyses

Elemental analyses for the muscovite and biotite sample were obtained using the

fully automated 3-spectrometer Geology and Geophysics Dept. Cameca microprobe

operated in a wavelength dispersive mode. Probe analyses were made using thin-sections

of the two samples. The minerals used as standards (and the elements for which they were

used) were Amelia albite (Na), Orthoclase Or-I (K), Smithsonian Anorthite (AI, Si,

Caj.Marjalahti Olivine (Mg) and the University of New Mexico Ilmenite (Ti, Mn, Fe).

A 15 KV filament voltage, lOnA beam current, defocused beam diameter of 10 urn

and a counting time of 10 seconds per element were used in all cases. The standards were

checked periodically for instrument drift. Raw data were adjusted on-line for detector dead

time and system background, and analyses were reduced using ZAF correction. Each

reported analyses represents the average of at least 5 spots per thin-section. Table 17 gives

the elemental composition of the muscovite and biotite samples. The compositions of the

phlogopite have been reported previously ( Matson et. al., 1986). Data for a representative

sample (BD3088) are given in Table 12 .

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