J . Phys. Chem. 1987, 91, 503-505 it is also consistent with the 20% conversion extent of Ce3+ into Ce4+ reported previously" under oscillatory conditions for the cerium-catalyzed BZ reaction. Thus this gradient is about M mm-' along the wavefront, once the full amplitude has been reached. This first experimental estimate of the order of magnitude of a concentration gradient associated to a trigger wave in an oscillatory medium provides a new criterium for evaluating the different available models. Finally, Figure 2c reveals a slight increase of the propagation speed as the distance increases from (1 1) Vidal, C.; Roux, J. C.; Rossi, A. J. Am. Chem. Soc. 1980,102, 1241.
503
the center. Indeed, for r < 0.5 mm no value greater than 90 pm s-' is obtained, whereas all values are above this threshold for r > 2 mm. The observed speeds (a100 pm s-I) are consistent with previous determinations made farther from the center.6 Unfortunately, the lack of accuracy prevents us from concluding that the asymptotic speed is reached 2 mm away from the center, as we would expect. Nor can these measurements support in a convincing manner any mathematical relationship between speed and distance. Further progress now requires an enhancement of the sensitivity of optical detection, hopefully making possible the use of higher magnifications.
Calorimetric Study of the Vitrified Liqr-'d Water to Cubic Ice Phase Transition Andreas Hallbrucker and Erwin Mayer* Institut fur Anorganische und Analytische Chemie, Universitat Innsbruck, A 6020 Innsbruck. Austria (Received: September 23, 1986; In Final Form: November 1 1 , 1986) The vitrified liquid water to cubic ice phase transition was investigated by differential scanning calorimetry and its enthalpy determined as -1.33 & 0.02 kJ mol-'. A second weak and broad exotherm extending from 125 to approximately 175 K might be due to enthalpy relaxation in the amorphous phase; it contributed 12 & 1% to the total heat release. A comparison with the literature data for H,O(as) suggests that less heat is evolved during the phase transition of vitrified liquid water than that of H20(as).
Introduction It has been debated recently whether or not the amorphous forms of water made from the liquid state and from water vapor have the same structure.'** The vapor deposited amorphous form, H20(as), has been known since 19353and investigated in great detail;2 the vitrified liquid, on the other hand, became available only r e ~ e n t l y . ~Early investigations of the latter were hampered by the presence of liquid c r y ~ m e d i u m . ~Liquid water can now also be vitrified by a new technique without liquid cryomedium, i.e., by rapid cooling of aqueous aerosol droplets on a cryoplate.68 This technique gives vitrified samples which are particularly suitable for further investigations of their properties. So far the infrared spectrum' and differential thermal analysis warm-up curves4 have been reported. In this letter a study of the phase transition from vitrified liquid water to cubic ice by differential scanning calorimetry, DSC, is reported and the enthalpy of transition compared with the literature values for H20(as). Experimental Section The vitrified liquid water samples were prepared as described in earlier Briefly aqueous aerosol droplets, made by an ultrasonic nebulizer, were transferred through a small opening together with nitrogen as carrier gas into a high-vacuum system, accelerated by supersonic flow, and deposited on a cryoplate. The only impurity was some codeposited water vapor. Continuous deposition of aerosol for 40-60 min gave a deposit 0.5-1 mm thick in the form of a disk. The amount of crystalline ice in the quenched samples was determined by X-ray diffraction by comparing the intensity of the peak from the 220 cubic ice reflection before and after devitrification. Optimally quenched samples, made as described in ref 7 and 8, were used which contained (1) Johari, G. P. Philos. Mag. 1977, 35, 1077. (2) Sceats, M. G.; Rice, S. A. In Treatise on Water; Franks, F., Ed.; Plenum: New York, 1982; Vol. 7, Chapter 2. (3) Burton, E. F.; Oliver, W. F. Proc. R. SOC.London A 1935, 153, 166. (4) Mayer, E. J . Microsc. (Oxford) 1985, 140, 3 and references therein. ( 5 ) Mayer, E.; Briiggeller, P. J. Phys. Chem. 1983, 87, 4744. (6) Mayer, E. J . Appl. Phys. 1985, 58, 663. (7) Mayer, E. J. Phys. Chem. 1985, 89, 3474. (8) Mayer, E. J . Phys. Chem. 1986, 90, 4455.
0022-3654/87/2091-0503$01.50/0
consistently 5 f 1% crystalline, mainly cubic, ice. The contribution of water vapor to this value of 5% ice was assessed in a separate experiment as described previously (ref 17 in ref 8) where water vapor from saturated nitrogen was deposited on a cryoplate without water droplets, but with otherwise identical conditions. This gave 30% crystalline, mainly cubic, ice. The value of 95% vitrified material from aqueous aerosol is therefore a lower limit with respect to the water droplets because the vapor in the aerosol will contribute significantly to the remaining 5% crystalline ice. A Perkin-Elmer DSC-4 instrument computerized with the TADS system was used. Water, cyclohexane, cyclopentane, and n-heptane (Merck, Uvasol quality) were used for calibration. Calorimetric accuracy of the low-temperature phase transitions was within specification, Le., f 1%. The transition temperatures were accurate to better than f l K. Curvature was eliminated with the S A Z function which subtracts during scanning a base line obtained with empty cells. We found the high-pressure cells offered for the DSC-4 instrument most suitable because the cells could be filled and closed under liquid nitrogen, and the comparatively high mass of the cells prevented excessive warming during rapid transfer from liquid nitrogen into the precooled instrument. The cells were filled with small chips of quenched sample which had been deposited either on a Cu cryoplate or on an X-ray low-temperature sample holder for simultaneous determination of the percentage of crystalline ice. Approximate sample weights were between 4 and 7 mg. The samples were scanned from 103 to 283 K with a heating rate of 10 K min-'. We were not able to determine the mass of water in the cells by weighing with sufficient accuracy: the cells which had been closed under liquid nitrogen leaked at room temperature and some of the water had vaporized even within the few seconds necessary for rapid transfer from the calorimeter to an automatic balance. Therefore we determined the mass of water from the heat of melting, using the value of 6.012 kJ mol-'. This amounts to a ratioing of the two heats of transition by their peak areas. This approach was tested with cyclohexane which has two crystalcrystal phase transitions, and gave values within specification. The above method for determining the mass of water presupposes that water vapor does not condense onto the cold cell during transfer to the precooled instrument because any additional ice 0 1987 American Chemical Society
Letters
504 The Journal of Physical Chemistry, Vol. 91, No. 3, 1987
Figure 1. (a) Thermogram of a vitrified liquid water sample, heating rate 10 K min-'. Broken lines indicate onset and minimum temperatures,vertical
bars limits of integration. (b) Ordinate fivefold expanded. TABLE I: Enthalpy Chnnge AH during the Transformation from Vitrified Liauid Water to Cubic Ice' ~~
TABLE Ik Comparison of Enthalpy Changes AH during the Transformation from the Amorphous Forms of Water to Cubic Ice
~
AH/kJ
T/K
sample 1 2 3
4 mean
onset 158.7 158.3 159.0 158.3 158.4 158.3 158.7 158.7 158.5 f 0.2
minimum 163.0 162.7 163.3 162.6 162.4 162.5 162.9 162.7 162.8 f 0.3
mol-'
-1.287 -1.289 -1.273 -1.251 -1.238 -1.257 -1.252 -1.273 -1.265
f
0.019
AH corrected for contamination by 5 f 1% ice = -1.33 f 0.02 kJ
mol-'. would lower the value of the heat of devitrification. For optimal conditions, i.e., rapid flushing of the glovebox with dry nitrogen before and during transfer and 3.5 bar pressure of dry helium for flushing the instrument head, we found condensation during transfer negligible: in a blank run with an empty cell, handled in the same way as filled cells, the heat effect in the melting temperature region of ice was 10.3% compared with filled cells. Evaporation of water was no problem when the samples were warmed only to 283 K and rapidly cooled again: this gave nearly identical melting peak areas. Therefore we are confident that we have made no error due to evaporation losses. Results It has been reported that vitrified liquid water devitrifies in two steps,4" the intensity of the high-temperature exotherm increasing with the relative amount of vitrified material. In the samples used in this work only one exotherm was observable due to further improvements in quenching conditions, with an indication of a weak shoulder on the low-temperature side in some scans. Figure la shows a thermogram with the exotherm from a vitrified liquid water to cubic ice phase transition. The two bars indicate the limits of integration for determining the peak area. The data are summarized in Table I. Samples 1 and 2 were prepared by aerosol deposition at 77 K, and samples 3 and 4 by deposition at 60 and 94 K, respectively. The amount of crystalline ice impurity, determined on the same samples, was 5 f 1%.
amorphous form vitrified liq water H,O(as) H,O(as) H2O(as) H,O(as) low-density amorphous ice (from ice I)
AH/kJ mol-' -1.33 f 0.02 -1.8 f 0.2 -1.64 -1.5 f 0.1 -0.9
-1.425
ref this letter 10
9 11 12 13
Therefore the mean enthalpy of transition corrected for contamination by ice is -1.33 f 0.02 kJ mol-'. The samples contained also some H20(as) which, according to the infrared ~ p e c t r u m , ~ was estimated to be 3%. The data in Table I were obtained with a heating rate of 10 K mi&. We have in addition investigated the influence of heating rate: with 5 K min-I, both onset and peak minimum temperatures were shifted 3-4 degrees to lower temperatures, and with 20 K min-I, by the same amount to higher temperatures. A further broad and weak exotherm, extending from 125 K to somewhere within the phase transition, was observed in all the thermograms. It is appearent more clearly in Figure 1b which is a fivefold expansion of Figure la. Figure l b contains in addition the trace from a second run, obtained after stopping the first scan at 180 K and cooling the sample again. A comparison of the two traces shows first that the weak exotherm is not an artifact caused by the instrument, and second that it is an irreversible effect. The peak maximum appears to be between 140 and 150 K. An enthalpy estimate of approximately -0.19 kJ mol-' is obtained by integration between 125 and 176 K (long bar) and by subtracting the value of the phase transition (short bar). This amounts to 13% of the total exothermic effect. The curve subtraction procedure was reproducible and gave 12 1% for seven samples; however, we feel that this value might be quite inaccurate and should be considered only as an upper limit because the crystallization exotherm might contribute to it. In a separate experiment we tried to ascertain that the weak exotherm is not, or only to a minor amount, caused by some of the quenched deposit transforming to cubic ice. Annealing of a sample at 148 K in the X-ray diffractometer gave after 1 min no indication for the formation of cubic ice; after 6 min only about 3% of the amorphous phase had devitrified. This annealing temperature is right in the middle of the broad exotherm and in case of a phase transition
*
Letters we would therefore have expected to see immediately the formation of cubic ice. In accordance with earlier reports4v5we observed no evidence for a reversible glass transition in the thermograms of vitrified liquid water samples, despite much higher sensitivity with the new instrument. Discussion Table I1 compiles the heats of transition from the various forms of amorphous solid water to cubic ice. For the H20(as) to cubic ice phase transition all the three more recent studies+" give higher values than that determined in this work. Only the first studyl2 gives a much lower value, possibly due to contamination of the sample by crystalline ice. The values for H,O(as) probably contain an exothermic contribution due to reduction of surface area. Such an effect must be much smaller in the vitrified liquid water samples because quenched droplets of approximately 3 pm diameter were transformed to cubic ice. In low-density amorphous ice prepared by "melting" ice 1 at 77 K by pressure and annealing of the high-density amorph," no reduction of surface area can occur. The structure factor of this material is similar to that of H20(as).l4 Its value of -1.425 kJ mol-' can thus be seen as a lower limit, free of surface area heat effects. Despite these uncertainties, the differences between the heats of transition of vitrified liquid water and H20(as) seem to be outside the experimental error.15 Our X-ray data indicate that the broad exotherm observable in Figure l b at lower temperatures is not caused by the phase transition to cubic ice. Therefore we suggest that it is due to enthalpy relaxation in the amorphous phase. A comparable transition has not been reported in the studies of HzO(as). There, for similar rates of heating, heat evolution due to surface area (9) Sugisaki, M.; Suga, H.; Seki, S. Bull. Chem. SOC.Jpn. 1968,41,2591. (10) Ghormley, J. A. J . Chem. Phys. 1968, 48, 503. (11) MacFarlane, D. R.; Angell, C. A. J . Phys. Chem. 1984, 88, 759. (12) Beaumont, R. H.; Chihara, H.; Morrison, J. A. J. Chem. Phys. 1961, 34, 1456. (13) Handa, Y. P.; Mishima, 0.;Whalley, E. J . Chem. Phys. 1986, 84, 2766. (14) Bosio, L.; Johari, G. P.; Teixeira, J. Phys. Rev. Lett. 1986, 56, 460. (15) We have desisted from a comparison of the thermal behavior of
vitrified liquid water with that of H20(as) reported recently," although the latter was investigated by DSC using similar conditions, first, because in that report no heating rate was given, and second, because the devitrification of H20(as) can be influenced strongly by the presence of oxygen or nitrogen, exotherms occurring in air between 155 and 200 K for a heating rate of 6 K min-' (Mayer, E.; Pletzer, R. J . Chem. Phys. 1984,80,2939), and the DSC study of H,O(as) in ref 11 had been performed at normal pressure in the presence of air.
The Journal of Physical Chemistry, Vol. 91, No. 3, 1987 505 reduction has been reported to cease between 133 and 143 K.I0 The low-temperature end of the exotherm a t 125 K might be artificial, caused by partial annealing during thermal equilibration of sample and cell at 103 K before the scan. It has been suggested by one of the reviewers as an alternative interpretation that the broad exotherm may be the result of the completion of a nucleation process from prenucleation clusters which form during quenching in the liquid and which then stabilize during reheating; these would be far too small to give detectable X-ray lines but would contribute to the total heat of crystallization. We cannot exclude this alternative but consider it unlikely because we had observed a similar exothermic effect even in concentrated aqueous solutions quenched the same way, for example, in a 50% ethylene glycol-water solution. In this system no heat release would be expected from the completion of a nucleation process because the rate of nucleation in a concentrated solution is much lower than in pure water.16 In our opinion this contribution does not settle the debate whether or not vitrified liquid water and H,O(as) are the same structurally. Experimental evidence for and against it has been reported from a comparison of electron diffractograms" and of infrared spectra.' We cannot exclude the possibility that the higher enthalpy values observed for the devitrification of H,O(as) might just be caused by additional exothermic heat effects due to reduction of surface area. A crucial experiment for a structural comparison, which is being planned, is to obtain high-quality diffraction data, extended to scattering at large k to see differences in short-range order, for both the vitrified liquid and H,O(as). The absence of an observable glass transition in vitrified liquid water parallels what has recently been reported in a DSC study of HzO(as).I1 Both are at variance with an extrapolated value of about 139 K for pure water obtained from concentrated binary solution data1*and with a similar TBvalue obtained by adiabatic calorimetry for H , O ( ~ S ) . ~We do not want to overinterpret this apparent similarity with the DSC study of H,O(as)" or the contradictions at present, but hope to obtain a better understanding from future studies of vitrified dilute aqueous solutions. Acknowledgment. Support by the "Fonds zur Forderung der wissenschaftlichen Forschung" of Austria is gratefully acknowledged. We are indebted to Prof. G. P. Johari for discussing the manuscript. (16) MacFarlane, D. R.; Kadiyala, R. K.; Angell, C. A. J . Chem. Phys. 1983, 79, 3921. (17) Dubochet, J.; Adrian, M.; Vogel, R. H. Cryoletters 1983, 4, 233. (18) Angell, C. A,; Sare, E. J. J . Chem. Phys. 1970, 52, 1058.
