J . Chem. SOC., Faraday Trans. I , 1985,81, 2421-2435

Solvation of Nitroxides'f

BY MARTYN C. R. SYMONS* AND ANTONIO S. PENA-NU%EZ

Department of Chemistry, The University, Leicester LE1 7RH

Received 10th December, 1984

The e.s.r. spectra of nitroxide radicals are widely used to study environmental effects, especially for labelled biomolecules. However, e.s.r. spectra are, in general, time-averaged records of the environmental constraints experienced by the > NO' species. This averaging often shows up as a contribution to the linewidths, but the hyperfine coupling, A(14N), and g values are time-averaged values.

We have used non-averaging spectroscopic techniques (infrared, Raman and electronic spectroscopies) to obtain further information about the structures of nitroxide solvates. Unfortunately, the environmental sensitivity of the N-0 stretching frequency proved to be so low that definitive results were difficult to obtain without careful computer analyses. Since the visible absorption band at ca. 23000 cm-l can be measured with great accuracy, it was used in the major part of this study.

Systematic variations in v,,, for the 23000 cm-l band for pure solvents followed a similar trend to that for the >C=O stretFhing frequency of acetone. This suggests that the mode of solvation for the>C=O and>N=O units is similar. For mixed solvents, in contrast to previous work on acetone systems, we were not generally able to resolve the experimental spectra into two or more components because the shift-width ratios were too small. However, by experimenting with a series of computer-constructed curves, we have established that skew bands comprising two Gaussian bands of very different intensities can be uniquely analysed to give the position of the weak band. Hence, we have shown that for water+aprotic-solvent systems nitroxides form three distinct solvates, whilst methanol + aprotic-solvent systems form only two. Since the intermediate band of the aqueous systems is close to that for the nitroxide in methanol, we postulate that water forms two hydrogen bonds, falling to one as the aprotic solvent is added, whereas methanol only forms a monosolvate. These conclusions are identical with those deduced for acetone solutions, using the C=O stretching band as a gauge. Detailed comparison shows that the nitroxides behave as slightly stronger bases than the ketones. These results are used to reinterpret certain e.s.r. results, and it is suggested that nitroxide spin labels used in the study of membranes are often mono- rather than di-hydrated.

The importance of e.s.r. studies of nitroxide radicals (R,NO) in the field of 'spin labelling' of biological molecules cannot be overestimated, as is evidenced by the extensive and growing literature on this subject.'. Most of these studies are concerned with measurements of T,, the correlation time for the probe, a technique which has recently been extended into the very slow tumbling regime by the use of saturation- transfer e.s.r. s p e c t r ~ s c o p y . ~ ~ ~ To a great extent this aspect of e.s.r. studies has overshadowed the other important parameter, namely the 14N hyperfine coupling constant, A(14N). This is sensitive to changes in environment, and can be used, for example, to decide if the )NO group is in an aqueous or a hydrocarbon environment. As we indicate below, it may be significant that in biological systems the A(14N) value frequently falls about half-way between these two limiting values.

We and others have frequently used e.s.r. spectroscopy to study solvation

t Taken as Solvation Spectra, Part 77.

242 1

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2422 SOLVATION OF NITROXIDES

p h e n ~ m e n a , ~ - l ~ in particular, we have used nitroxide radicals as probes in the study of mixed-solvent systems13 and of electrolyte s01utions.l~ The work of Jolicoeur and Friedman,l59 l6 mainly centred on linewidth effects for solutions of TEMPO (2,2,6,6- tetramethylpiperidine-N-oxide and OTEMPO (2,2,6,6-tetramethyl-4-piperidone-N- oxide), was also concerned with solvation effects in pure and mixed-solvent systems. There are marked interpretative differences in these two sets of studies, arising largely because we focused attention on changes in A(14N) and the way in which linewidths depended thereon, rather than relying solely on width variations. In the present work we call upon our own e.s.r. resultsL3 for comparison with the new data.

In other studies of probe molecules in pure and mixed-solvent systems we have shown how useful vibrational spectroscopy can be in the study of mixed-solvent systems. The main reason for this is that 'static' structures are detected, since the rate of interchange between different solvate structures is relatively slow on the infrared timescale. This is in marked contrast to the situation for n.m.r. studies, which usually give narrow, time-averaged lines, and even for e.s.r. studies, although in these cases line-broadening due to interchange between different solvate structures may make significant contributions to the overall widths.13

Thus, magnetic-resonance studies give time-averaged data which are nevertheless very precise, whereas vibrational or electronic spectroscopic studies give non-averaged data; however, since in the latter case the features are broad, the information is much less precise. We have recently shown how careful analysis of the infrared results for solutions of triethylphosphine oxide, Et3P0, gives structural data which can be used to obtain a quantitative reconstruction of the 31P resonance shifts for mixed-solvent systems. l7

As stressed above, the major limitation of infrared and ultraviolet spectroscopic studies of solvation effects is that the significant features may be very broad. If their widths are large in comparison with solvent-induced shifts, then analysis is difficult, and it may only be possible to study shifts rather than changes in component features.'** l9

The aim of the present study was to measure changes in vibrational and electronic spectral features for nitroxides in the hope that one could be found which gives information about separate solvates in mixed-solvent systems, and hence reinterpret some of the e.s.r. results in the light of such information.

EXPERIMENTAL Solvents of the highest grades available were purified and dried by standard procedures.

Water was purified using a Millipore ' Milli-Q' system after deionization. When necessary, solutions were prepared in a dry box.

Visible spectra were measured with a Perkin-Elmer 340 spectrometer. For most studies, 10 mm stoppered silica cells were used. Temperatures were controlled to fO.1 "C. Infrared spectra were measured with a Perkin-Elmer 580 spectrometer, using demountable cells, usually with calcium fluoride windows and a pathlength of 0.25 mm. E.s.r. spectra were measured with a Varian El09 spectrometer using flat silica cells. Samples were thoroughly deoxygenated before use.

Resonance Raman spectra were measured at York University using a Jobin Yvon Ramanov RG2 spectrometer linked to a Nicolet 1074 data system, by courtesy of Prof. R. E. Hester. A Spectra Physics 170 krypton ion laser was used to give excitation at 413.1 nm, and all spectra were obtained using the spinning-cell technique.

Digitized spectra were processed using a Gaussian curve-analysis program and DEGCT42 graphics display linked to a PDP 1 1 /45 minicomputer.

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M. c . R. SYMONS AND A. s. P E N A - N ~ E Z 2423

RESULTS A N D DISCUSSION

CHOICE OF ABSORPTION BANDS

Our first aim was to select the best absorption feature for study. In previous work on acetone,20v21 to which we refer extensively herein because of the structural similarities between ketones and nitroxides, the carbonyl stretching mode was most satisfactory and we had hoped that this would also be the case for the N-0 stretching feature in the infrared spectrum of nitroxide systems. However, this feature, which occurs at ca. 1340 cm-l for di-t-butylnitroxide (DTBN), is relatively weak, and is close to other more intense features.22 Furthermore, it proved to be insensitive to changes in environment, in contrast to the C=O stretch of 21 In our experience such insensitivity does not necessarily reflect weak solvent interactions; in this case we suggest that it is of structural origin, possibly arising because of the presence of an antibonding n-electron with resulting competing effects from solvent perturbations. Nevertheless, several solvent systems were studied, and the results tie in well with our more extensive visible-absorption studies described below.

We also investigated the Raman spectra of a range of solutions, hoping that those would give a better defined and possibly more sensitive feature. To improve the sensitivity of the N-0 stretching feature, resonance Raman spectroscopy was used, since this selectively enhances features from the NO group, which is the optical chromophore. Certainly, the N-0 shift dominated the spectra, but nevertheless the feature was weak and reproducibility was poor. As with the infrared studies the shifts were small, but our limited studies gave results in accord with those described below.

In all subsequent work we used the visible (n + n*) band for the nitroxides to monitor solvation effects. This enabled us to use dilute solutions and to obtain the reproducibility needed for careful computer analysis.

PURE-SOLVENT SYSTEMS

Band maxima for nitroxides in a range of solvents, given in table 1, are displayed as a function of various other parameters in fig. 1. The linearity of the plot in fig. 1 (a) shows that the factors controlling the 14N hyperfine coupling are directly linked to those controlling the band maxima. The only points which deviate significantly from this correlation are those for t-butyl alcohol and dimethyl sulphoxide. We have no explanation for the latter, but discuss the former below. For the remainder, this correlation suggests that only one major type of solvate is present in these solutions. This accords with the symmetrical Gaussian form of the absorption bands. As we argue below, had there been other components of low intensity, these bands would not be symmetrical.

In fig. 1 (b) we compare our results with the band maxima for the C=O stretch of acetone in the same solvents. In this case two correlating lines are required, one for the aprotic and the other for the protic solvents, both sharing the inert solvent value. In our experience of a range of solvent probes this separation is normal. It occurs because the mode of interaction in protic solvents involves quite specific hydrogen bonding, whilst that in aprotic solvents involves an averaged dipolar interaction which, in general, has a different perturbing effect on the chromophore. Indeed, in some cases the correlating curves have opposite slopes. Even so, the scatter for the aprotic solvents is large, especially for hexamethylphosphoramide (HMPA) and cyanomethane (MeCN). For HMPA the nitroxide results suggest a far lower polarity than the acetone results. We suggest this reflects the large steric barrier afforded by the Me&- and Me2N- groups, which precludes the formation of an intimate dipolar interaction for the nitroxide.

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2424

17.5

17.0-

SOLVATION OF NITROXIDES

-

Table 1. DTBN information

&? 165-

v

16.0

peak position (25 "C) key solvent /cm-'

-

a b

d e f g h

C

1

j k 1

m n

P 9 r S t U

0

acetic anhydride toluene CCl, DMF DMA Et,NH HMPA acetone DMSO NMF ButOH MeOH

MeCN EtOH IPA cyclohexane dodecane ethanediol formamide NMA (30.6 "C)

H2O

22 173 21 786 21 786 22 075 22 002 21 882 21 882 22 026 22 099 22 523 22 523 22 83 1 23 753 22 222 22 676 22 573 21 622 21 622 22 962 22 962 22321

1

I 1 1

2 3500 23000 22500 22000 15.0'

vm&lx/cm-'

Fig. 1. Correlations between v,,, for di-t-butylnitroxide (DTBN) in a range of solvents and (a) the 14N hyperfine coupling, (b) infrared results [v(C=O)] for acetone, (c) the solvent acceptor number and (d ) as for (c), but using numbers corresponding to the dihydrate (a) and the monohydrate (p). (Key: see table 1.) [Data from this work and from B. R. Knauer and J. J. Napier, J. Am. Chem. SOC., 1976, 98, 15 for A(14N) values of DTBN in various

solvents.]

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24000

2300( - I

E -2 \ X

A

2200(

2100(

M. c. R. SYMONS AND A. s. P E N A - N ~ E Z 2425

m

1 I 1 I I

1700 1710 1720

vmax(C=O)/cm-l

2 4 0 0 0 c /

/ m

I

0 /*

/;''

/

0 10 20 30 40 50 60 acceptor number of solvent

Fig. 1 (b-d). For legend see opposite.

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2426 SOLVATION OF NITROXIDES

The correlating line for protic solvents [fig. 1 (b)] is good, with the exception of the ethanediol data. In particular, the shift from the inert-solvent value for methanolic solutions (Av = 1209 cm-l) is close to half that for aqueous solutions (Av = 21 3 1 cm-l). For both probes this could mean either that methanol forms weaker hydrogen bonds or that the bond strengths are comparable but that water forms twice as many as methanol (of course both factors could apply). We favour the latter extreme because we do not expect the B - - .HOMe bonds to differ greatly in strength from the B . - .HOH bonds. This is supported, for example, by our results for the solvates X- . HOMe and X- * HOH (where X is a halogen) in te t rachl~romethane.~~ For acetone, previous mixed-solvent studies revealed the participation of two major species in methanol + aprotic-solvent systems, but required three spectroscopically distinct species for water + aprotic-solvent 21 These results are under- standable if we postulate that most acetone molecules form two hydrogen bonds in dilute aqueous solution, but only one in methanoL21 Reasons for this remarkable contrast are discussed below. The good correlation between the protic-solvent results for DTBN and acetone strongly suggest that a similar solvation pattern applies in this case.

for the display. As we have come to expect,22 the linear correlation fails in that the key results for aqueous solutions are not accommodated. The aprotic-solvent data correlate well with those for the alcohols, but we suggest that this is fortuitous. A solvent acceptor number is the 31P resonance shift for triethylphosphine oxide (Et,PO) in that solvent.24 These should only correlate with other parameters for other probes if the solvation is comparable. Previous extensive infrared and 31P n.m.r. studies of Et,PO have shown, at least to our satisfaction, that the normal solvate in methanol contains two hydrogen-bonded solvent molecules, whilst that in water has three.18 In order to compare results with those for acetone or DTBN, we suggest that the 31P shift for the disolvate be used to correlate the water solvate data, and that for the monosolvate be used for the alcohol data.22 When this is done, the alternative correlation, (a), in fig. I(c) is obtained. In this case the protic-solvent values are now well accommodated, but by a line well separated from that for the aprotic media. We suggest that, used in this way, acceptor numbers may be useful, but one should not expect to find a single correlation for both types of solvent, nor will the protic-solvent data be properly accommodated unless allowance is made for variable solvation numbers.

In fig. 1 ( c ) we use the solvent acceptor

METHANOL + APROTIC SOLVENTS

Shifts in A(14N) and in vmSx as a function of mole fraction of cosolvent (MeCN or DMSO) are shown in fig. 2(a) and (b). These trends make an interesting contrast,

Fig. 2. Methanol + aprotic-solvent systems. (a) Trends in 14N hyperfine coupling constants. The smooth curves are the e.s.r. results and the points are from our analysis described in the text. (b) vmax trends. ( c ) Optical spectra for DTBN in MeOH+DMSO mixed solvents.

Key: solution MeOH mole fraction DMSO mole fraction

A 1 .o B 0.803 C 0.539 D 0.236 E 0.0

0.0 0.197 0.46 1 0.764 1 .o

~ ~~

( d ) Trends in the intensities of the two bands needed to simulate the spectra: (i) MeOH + DMSO, (ii) MeOH+MeCN and (iii) acetone in MeOH+MeCN [taken from ref. (21)].

(a = protic solvents; p = aprotic solvents.)

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M. c. R. SYMONS AND A. s. PENA-NUGEZ 2421

mole fraction of base

I I I I I 600 500 600

A/nm

100.

50

0. DMSO (mole fraction)

i loo*,.

h ‘a,. i P)

250- -e 2 m ./. \ .’ .’ 0 ./

MeCN (mole fraction)

(ii)

I I I I

0 0.2 0 .6 1.0

mole fraction of cosolvent

Fig. 2. For legend see opposite.

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2428 SOLVATION OF NITROXIDES

especially for the DMSO systems, which appear to give different information. In fact, as we show below, the e.s.r. trends are meaningfL1, but the vmax trends are misleading and have no direct significance.

For the optical studies, both systems revealed precise isosbestic points, as shown in fig. 2(c ) for the DMSO system. This strongly supports our contention that only two well defined species are involved, the monomethanol solvate and the non- hydrogen-bonded molecules. Curve analysis, using only the two pure-solvent bands, nicely reproduces the mixed-solvent bands. This gives us the proportions of the two species present in any one solution, their trends as a function of the solvent composition being shown in fig. 2(d) [(i) DMSO and (ii) MeCN]. These quantitative results confirm our qualitative expectations. Thus for the strong base, DMSO, the equilibrium

R2NO- . -HOMe+BeR,NO+B- . .HOMe ( 1 )

lies to the right, whilst for the weaker base MeCN it lies to the left. In fig. 2(d ) (iii) we show the comparable results for acetone probe2' in MeOH + MeCN mixtures. They show clearly that the nitroxide is a stronger base than acetone.

Trends shown in fig. 2(d ) show how one can be badly misled by plots involving shifts in vmax [fig. 2(b)]. The latter will be increasingly meaningful as the number of species contributing to the shifts increases. Thus had there been, say, five distinct species involved, the vmax shifts in fig. 2(b) would hake closely resembled the e.s.r. shifts shown in fig. 2(a). Only under such conditions are vmax shifts likely to be anything but misleading.

These analyses can be used to predict the e.s.r. shift data shown in fig. 2(a). The smooth curves represent the experimental trends and the points are taken from the optical data [i.e. from fig. 2(d)]. The agreement is good. This supports our contention that only two structurally significant species are present.

WATER + APROTIC SOLVENTS

As for the methanol systems, the trends in A(14N) and in v,,, are given in Fig. 3(a) and (b). However, for these systems, if our comparison with acetone is valid, we need three curves in the reconstructions. We have obtained results from mixed-solvent systems rich in water and rich in MeCN which strongly support this contention [fig. 3(c)]. As MeCN is added to aqueous DTBN, the absorption-band intensity falls and the bond becomes skewed on the low-frequency side. Attempts to reproduce this asymmetry using an extra weak band in the MeCN region (22220 cm-l) failed. However, an extra band at ca. 22300 cm-l gives a satisfactory fit. Similarly, when water is added to a solution of DTBN in MeCN the band becomes skewed on the high-frequency side, but cannot be reproduced by adding an extra feature at the pure- water value. Once again, a good fit is obtained on the addition of the same intermediate band.

These results are entirely consistent with our conclusions drawn above. The intermediate band is assigned to the monohydrate and falls close to the band for methanolic solutions. This conclusion receives further support from the set of curves in fig. 3(d). This shows two regions containing reasonable isosbestic points, one for water-rich solutions and one for MeCN-rich solutions.

Computer analysis of these spectra using three bands leads to the trends in relative concentrations shown in fig. 3 (e) for the water + MeCN and water + DMSO systems. Again these trends are similar to those obtained more directly for acetone.22 They suggest that the monohydrate dominates in the 0.35-0.65 mole fraction (MeCN) region. The disolvate is lost rapidly despite the low basic strength of MeCN. This arises because the two hydrogen bonds in the dihydrate are much weaker than the single

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M. c. R . SYMONS AND A. s. PENA-NU~EZ 2429

h

v c1

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2430 SOLVATION OF NITROXIDES

0 N

0 W

0 0

h

Y .- .-

0 0 0

5 \ <

x 3

D

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M. c. R. SYMONS AND A. s. PENA-NU~~EZ 243 1

monohydrate bond. As stressed elsewhere,22 we cannot compare the behaviour of the monohydrate with the methanol system [fig. 2(d)] directly because each water molecule provides two binding sites. However, in a qualitative sense this explains why the addition of water to MeCN solutions causes a more rapid conversion to the monosolvate than the addition of methanol.

The results shown in fig. 3(e) can be used to reproduce the e.s.r. trends shown in fig. 3 (a) . This has been done for H 2 0 + MeCN and H 2 0 + DMSO systems. The points in fig. 3(a) are actually the values predicted from the optical spectra and the smooth curves give the experimental hyperfine splittings. Once again, the good agreement strongly supports our analysis.

WATER + ALCOHOL SYSTEMS

The e.s.r. results and spectral shifts are shown in fig. 4(a) and (b), and a typical set of spectra are given in fig. 4(c) for the water+methanol system. Curve analysis was achieved using two bands, one for the dihydrate and the other for the monohydrate or the R 2 N 0 - - HOR species. Trends in the intensities of these bands are indicated in fig. 4(d). For H 2 0 + MeOH there is a remarkably steady change from the disolvate to the monosolvate. This reflects the similarity between these two solvents and suggests that their mixtures move smoothly from being water-like to being met hanol-like.

Analysis of the H 2 0 + t-butyl alcohol system makes a remarkable contrast. As alcohol is added, the trend begins in an identical manner to that for methanol, but then there is a catastrophic switch to favour the monohydrate or t-butyl alcohol solvate in the 0.1 mole fraction region [fig. 4(d) ] . This change suggests that the nitroxide moves into a different environment which makes disolvation impossible but which favours monosolvation. We suggest that this new environment is some form of cluster in which the nitroxide directly participates. It is clear that in regions of incipient phase separation such clusters occur, and we have presented ultrasonic relaxation25 and other results26 which favour such a concept for H 2 0 + ButOH systems. More recent results suggest some sort of microemulsion formation in this solvent 28

Fig. 3. Water + aprotic-solvent systems. (a) Trends in 14N hyperfine coupling. The curves are the e.s.r. results and the points are from our analysis described in the text: x , DMSO and 0, MeCN. (b) v,,, trends. (c) Skew band analysis for the optical spectrum of DTBN in water+MeCN mixtures. Note that the baselines for the two spectra have been displaced for clarity. The band labelled ‘skew’ is the computed difference curve which is responsible for the skewed nature of the H,O+MeCN curve, and this has been enhanced by a factor of 100 in curve a. ( d ) Experimental optical spectrum for the H,O+MeCN solvent system.

Key: solution H,O mole fraction MeCN mole fraction

1 .o 0.943 0.871 0.744 0.492 0.244 0.132 0.0

0.0 0.057 0.129 0.256 0.508 0.756 0.868 1 .o

( e ) Trends in the intensities of the three bands needed to simulate the spectra: (i) H,O + MeCN and (ii) H,O+DMSO.

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2432 SOLVATION OF NITROXIDES

ROH (mole fraction)

2380C

2310( 4

I

E 1

2300C

22600 1 I

0.6 0.8 1 .o ROH (mole fraction)

I I I 1 I I

Loo 500 6 00 h/nm

Fig. 4(a-c). For legend see opposite.

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M. c . R. SYMONS AND A. s. PENA-NUSEZ 243 3

0 0 . 4 0.8 ROH (mole fraction)

Fig. 4. Water+alcohol systems. (a) Trends in I4N hyperfine coupling: the full curves are the experimental data and the points are derived from the optical data-; A, MeOH and A, ButOH. (6) v,,, trends; 0, MeOH; a, EtOH; A, IPA and 0, ButOH. ( c ) Optical spectra for the H,O + MeOH solvent system (that for the H,O + Me,COH system was comparable).

Key: solution H,O mole fraction MeOH mole fraction

A 1 .o B 0.734 C 0.49 1 D 0.0

0.0 0.266 0.509 1 .o

( d ) Trends in the intensities of the two bands needed to simulate the optical spectra: (i) H,O + MeOH and (ii) H,O + Me,COH.

As before, we have used the results shown in fig. 4 ( d ) obtained from the deconvoluted optical spectra to reproduce the changes in A(14N). The agreement is again good, as can be seen from the points in fig. 4(a).

INFRARED AND RESONANCE RAMAN RESULTS

Some data from the infrared and Raman measurements are given in table 2. The Raman shifts are slightly greater than the infrared shifts, but are still too small for mixed-solvent work. As stressed above, these shifts are certainly not small because of weak interactions. Indeed, these are clearly stronger than for acetone, which displays large infrared shifts for the C=O group.2o We conclude that the small shifts reflect the different electronic structures, and that vibrational spectroscopy is of little use in the study of nitroxide solvation.

NITROXIDES IN MICELLES AND MEMBRANES

As stressed above, the 14N hyperfine coupling constants for nitroxide spin labels and for nitroxide spin traps act as a probe of the environment. It is frequently found that A(14N) is close to the value for methanol in such systems, i.e. ca. midway between the value for water and those for inert solvents. Various reasons for this have been

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2434 SOLVATION OF NITROXIDES

Table 2. Resonance Raman and infrared v(N0) (in cm-I) of DTBN in various solvents

Raman infrared

CD,OD 1342 MeCN 1340.0 CD,CN 1340.5 H*O 1337.5 C6Hl2 1347.0 C6Hl2 1342.0 THF 1337.0 CD,OD 1343.5 DMSO 1 343 .O H P 1335.0

proposed, including the concept of rapid equilibrium between the two environments. Two examples come from spin trapping in micellar systems. Thus Janzen and C o ~ l t e r ~ ~ state that ‘the environment probed by the nitroxyl function is still very polar and (is) approximately that of pure methanol’. Walter et af. reached a similar conclusion in related studies.30 We suggest that the species under observation in many of these studies is the monohydrate, which, as we have established, closely resembles the monoalcoholate.

Finally, we stress that, in contrast to our infrared and magnetic resonance studies, measurements using electronic spectra must give rise to net shifts caused by solvation effects on ground and excited states. In terms of our present aims, which effect dominates is of no particular significance. This is because distinct species are detectable, and these cannot have been formed during the process of light absorption. Thus they must already be present in the system, which is all we are concerned about.

We thank Dr Graham Eaton for extensive help and advice throughout this study. We are also grateful to Prof. R. E. Hester and Dr R. B. Girling for the use of their Raman spectrometer and experimental assistance.

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