Faraday Symp. Chem. SOC.,1982 17,31-40 'Infrared and Nuclear Magnetic Resonance Studies Pertaining to the Cage Model for Solutions of Acetone in Water* BY MARTYN C. R. SYMONS R. EATON AND GRAHAM Department of Chemistry The University Leicester LEI 7RH Received 3rd September 1982 Solutions of acetone in water show a single C=O stretching mode in the infrared (1697 cm-I), which is assigned to acetone hydrogen bonded to two water molecules. On cooling below 0 "Cnew features appear which are characteristic of concentrated solutions of acetone formed by growth of ice crystals. These bands are assigned to acetone forming single hydrogen bonds (1708 cm-') and no hydrogen bonds (1715 cm-l). The latter band is characteristic of bulk acetone. On standing at cu. -30 "C,all these features decay and ultimately only a single narrow band is detectable at 1722 cm-I which is characteristic of isolated acetone molecules in completely inert solvents.This change is clearly a result of clathrate formation. These results demonstrate unequivocally that acetone in liquid water is solvated in a manner that differs completely from cage "solvation". This remarkable difference between solid and fluid systems is interpreted in terms of the presence of large concentrations of free OH groups in liquid water. The concept of hydrophobic bonding is probably unique to aqueous systems. It is manifested in a gross manner in such systems as micelles and membranes and reflects the fact that every water molecule wants to form four hydrogen bonds water thus tending to reject molecules with which it cannot form bonds.For molecules that have regions that can form bonds and others that cannot there are a number of compromises micelles being a typical and satisfying example. Another related concept that is often invoked to explain the behaviour of water- rich aqueous systems is that of incipient clathrate cage formation. Water has a remarkable propensity to form fully hydrogen-bonded cages around inert solutes in the solid phase and although such cages are often nearly spherical they are a remark- able variety of more complex shapes in water's repertoire. An as yet unproven hypothesis invoked to explain a multitude of phenomena is that such cages are also important in the liquid phase at least at low temperatures.In our own studies such incipient cage formation was invoked to explain the appearance of intense ultrasonic relaxation effects in water + t-butyl alcohol systems,' proton resonance shifts for aqueous tetra-alkylammonium salts,2 asymmetry in the 6.s.r. spectrum for m-dinitro- benzene anions in aqueous sol~tions,~ and linewidths in the e.s.r. spectra for di-t- butylnitroxide in aqueous system~.~ Assuming for the moment that incipient cage effects may sometimes be significant in liquid-phase systems we need to enquire about possible links between these pheno- mena and hydrophobic effects. These are not obvious since in the solid clathrate compounds there is generally a filigree of water molecules separating each solute molecule thus even in these concentrated systems there is no tendency towards dimerisation or aggregation.This is not always true. For example two (C,H,),S+ * Taken as Solvation Spectra Part 73. CAGE MODEL FOR ACETONE SOLUTIONS cations share a single cage in the clathrate (C4H9)3S+F-8H20.5Thus cage-sharing may only be important when there are not enough water molecules to form mono- cages. Indeed the absence of ultrasonic relaxation behaviour in the 0-0.03 mole fraction range for t-butyl alcohol in water was taken as evidence that cage-sharing only set in at mole fractions > 0.03.' If this is correct we can argue that in systems favouring clathrate cage formation direct " hydrophobic bonding " will be avoided when the ratio H20 solute molecules is greater than that for the clathrate but increasingly favoured for smaller ratios.However another factor needs to be considered.'j In solid clathrates not only are the cages formed so as to give nearly perfect hydrogen bonding for the water molecules defining each cage but all cages are perfectly linked together. (By "per-fect '' we mean four strong nearly linear hydrogen bonds per water molecule.) Hence we can visualise a single cage as acting as a template for adjacent cages in the sense that a significant proportion of a second cage has already been constructed. This is even more the case for the construction of subsequent cages on a double-cage unit. These considerations lead to the tentative postulate that "cages breed cages ". If this were true then it would also lead to a kind of dimerisation or even oligomeris- ation of the solute.The two cases can be compared with the concept of contact and solvent-shared ion pairse7 If these two types of structure can be linked under the genetic term "hydrophobic bonding " perhaps they should be similarly differentiated as contact and solvent-shared units. The aim of the present work was to use infrared spectroscopy to probe similarities or differences between the state of a solute in liquid water and in its solid clathrate compound. We selected acetone as one example,' because of our extensive studies on the use of acetone as a spectroscopic probe of pure and mixed solvent ~ystems.~ In one of our previous spectroscopic excursions into systems that might be in- fluenced by water cages we suggested that the proton resonance for water molecules at low temperatures was shifted to low fields by R4N+ions the low-field contribution increasing as the size of R increased.2 This seemed to correlate with the well known ability for R,N+ salts to form clathrate hydrates.This trend was opposite to that previously detected 'OJ' at higher temperatures and we confirmed that the trend was indeed reversed on heating. This marked temperature effect accords with the fact that cage formation requires considerable organisation of many water molecules but the reversal of the trend seems to imply more than a loss of specific encagement at higher temperatures. It is relevant that shifts caused by R4N+ions in methanol are negligibly sma11.12 There are serious problems associated with this work,2 since we neglected to allow for changes induced in water by the halide ions.13 We postulate that there is a relatively high concentration of water molecules forming three hydrogen bonds in liquid water and we describe these broken bonds as (OH)f,ee and (LP)f,ee,where LP represents the " lone pair " of electrons not involved in hydrogen bonding.Despite the fact that this postulate is not generally accepted we use it herein without further ju~tification.'~ Although (LP)freegroups have never been detected spectroscopically it is usually considered that groups make a considerable contribution to the 0-H stretch overtone band^.'^-'^ When tetra-alkylammonium salts are added to water there is a sharp decrease in the overtone band at 7120 cm-' assigned to groups.This fall has quite reasonably been assigned to the effect of the R4N+ions on water since alkali halide salts induce much smaller changes in this The marked loss of (OH)freegroups was therefore taken as evidence for the " structure-making " ability of R4N+ions. This seemed unlikely to us since there is no major temperature effect on this behaviour and hence no correlation with the proton reson- M. C. R. SYMONS AND G. R. EATON ance results. In our view l8 the major part of this decrease in (OH)free groups is due to the reaction hal-+ n(OH)f,ee-+ hal(H0); (1) when n is the primary solvation number of the halide ion hal-. Given that R4N+ ions do not interact significantly with (LP)r,eegroups as is the case for methanolic solutions reaction (1) must cause a rapid loss of (OH)freegroups.For alkali-metal salts the metal ions generally interact with about the same number (n) of (LP)r,ee groups so that the net effect on equilibrium (1) is small. We expect that there is a small positive or negative contribution from R,N+ ions depending on the temperature but the experimental errors are unfortunately large and in our view such changes are not yet established. If these ideas are correct then the interpretation of the n.m.r. shifts previously proposed must be revised. This is in hand. In our view the e.s.r. results for rn-dinitrobenzene anions in water are more com- pelling and may constitute one of the best pieces of evidence in favour of significant structure around non-bonded groups in liquid water at low temperature^.^ Theory and e.s.r.studies in aprotic solvents require that m-dinitrobenzene anions be sym- metrical with two equivalent nitro group^.'^ However on forming ion-pairs with alkali-metal cations there is a dramatic loss of symmetry the spin-density on one NO2 group increasing by a factor of ca. 2 with that on the other falling to ca. 0.20 This unique modification can be understood in terms of the diagram shown in fig. 1.21*22 FIG.1.-(u) Combination of the ground-state SOMO (wScl,) for m-dinitrobenzene anions and a low-lying excited state ('yscz,) to give two asymmetric levels (yAs(l)+ I+Y*~(~)). These can be thought of as being primarily localised on one or other nitro group the asymmetric wavefunction being stabilised by asymmetric hydration as indicated 'in (6).The important result is that the e.s.r. spectrum in water is that of the asymmetric ion. From the linewidths governed by the transfer of charge and spin-density from one NO2group to the other a residence time z of 4.5 x s was calculated at 0 "C. This fell to ca. 0.8 x s at 25 "Cand rapidly decreased on heating. The lifetime CAGE MODEL FOR ACETONE SOLUTIONS increased by a factor of ca. 2 on adding 0.02 mole fraction of t-butyl alcohol but ul- timately fell to ca. s in the pure alcohol. Times in this region were also obtained for solutions in methanol. Although heating decreased z for aqueous solutions the 14N hyperfine coupling was unchanged showing that the extent of hydrogen bonding at the (NO2)- unit is independent of temperature.The only explanation that we can discover which can accommodate these results is that in protic solvents hydrogen bonding at one NO group pulls the negative charge onto this and away from the other NO2 group. This descent in symmetry tips the wavefunction into one of the two asymmetric forms (fig. 1). This perturbation will switch when some hydrogen bonding builds up fortuitously on the neutral NO2group and is concomitantly shed at the (NO2)- group so that they become effectively equivalent. Why should this be ca. times slower in water than in methanol? Our suggestion is that there is a cage around the neutral portion of the anion (fig. 1) and this needs to be destroyed as solvation builds up at neutral NO2.This very extensive reorganisation is not required for alcoholic solutions. The reinforcing effect of t-butyl alcohol in the 0-0.02 mole-fraction range may be evidence in favour of solvent-shared hydrophobic interactions. Another observation that led us to postulate incipient cage formation in aqueous solution is the marked downfield shift observed in the OH proton resonance for aqueous t-butyl alcohol. This shift which is much greater than that caused by methanol was initially interpreted by ourselves and by Hertz and coworkers in terms of incipient cage f~rmation.~~-~~ However Covington and Newman 26 concluded that the average shift observed was caused by a large downfield shift of the Me,COH protons and the water proton shift was unexceptional.Since then we have been able to obtain conditions of slow exchange and hence have observed both OH protons ~eparately.~' The results confirm our original interpretation it is indeed the water protons that display an anomalously large downfield shift in the usual 0-0.03 mole-fraction range. The only reasonable interpretation of this excess shift is that the cage- forming water molecules experience slightly stronger bonding than those in bulk water. The effect may also shift equilibrium (1) slightly to the left. Finally our studies of aqueous di-t-butyl nitroxide radicals in aqueous solutions In fact this claim was also made very give strong support to the cage the~ry.~.~~ strongly by Jolicoeur and Friedman 29 for similar systems but it seems to us that unfortunately their interpretation is in error.They interpreted their results in terms of a marked increase in zJ the spin-rotation correlation constant on cooling. This would mean that the R2N0molecules were gaining in rotational freedom on cooling which is most unexpected. They claim that this implies a gain of encaged R2N0 molecules which are not anchored to the cages and therefore are free to rotate. We showed that the increase in linewidth on cooling is due to an increase in zc the rotational correlation rather than in zJ. Furthermore the large magnitude of the hyperfine splitting A(I4N) which is indicative of strong hydrogen bonding does not alter in the significant temperature range (0-25 "C),and hence there cannot be any major trend towards non-hydrogen-bonded encaged molecules.Nevertheless we agree that the results imply cage formation although these must be anchored guests. Thus the cages confer unusual lack of freedom. On the addition of t-butyl alcohol zc increases still further in the 0-0.03 mole fraction range. We postulate that both additives are solvated in an elaborate manner at ca. 0 "C with hydrogen bonds linking both guests to partial clathrate cages and with cages reinforcing cages. On further addition of the alcohol zJ does begin to increase showing that the R2N0 molecules are gaining their freedom. Furthermore there is then a dramatic fall in A(14N) confirming that the N-0 HO bonds are lost. Evidently at this stage if cages M. C. R. SYMONS AND G.R. EATON are important the alcohol has tipped the equilibrium to the " free " stage envisaged by Jolicoeur and F~iedman.~~ EXPERIMENTAL Acetone was purified by standard procedures directly prior to use. D20 (99.8 atom %) was used instead of H20 because of the strong underlying absorption in the region of interest of H20. Infrared spectra were recorded on a Perkin-Elmer 580 spectrometer using demountable cells with Irtran windows and 0.025 mm path-length. Temperature was maintained using a Specac controller and thermostatted cell holder. ''0 n.m.r. spectra were recorded on a Bruker WM 400 Fourier-transform spectrometer and n.m.r. spectra on a Jeol FX 60 Fourier-transform spectrometer. Two methods of forming the clathrate were found to be successful (a) A suitable sample of acetone in D20(0.1 mole fraction) was cooled to -30 "C and held at this temperature for several hours the clathrate band was slowly seen to "grow" the effect being complete after cu.3 h. (6) The sample [as in (a) above] was cooled to below -100 "C and then allowed to warm up slowly at 2 "C min-I. The clathrate band was first observed at ca. -40 "C being fully formed at -38 "C. It disappeared rapidly at -19 "C. Method (6) was utilised for the work where alcohol was added to the system. RESULTS AND DISCUSSION INFRARED SPECTROSCOPY The key result of our infrared studies of the C=O stretching band for acetone in fluid water and in its clathrate is that there is a major shift together with a marked change in band shape (fig.2). We interpret this in terms of acetone forming two hydrogen bonds to OH groups in fluid water but no hydrogen bonds in the clathrate but this statement requires justification. wavenumber/cm-FIG.2.-Infrared spectra (C=O stretch) for acetone in water (0.05 mole fraction acetone) (a) at 0 "C (b)at -30 "Con storage for 7 h. Band maxima for acetone in (i) clathrate cage (ii) hexane (iii) bulk acetone (iv) concentrated and (v) dilute aqueous solutions are given as vertical bars. CAGE MODEL FOR ACETONE SOLUTIONS The clathrate band (1722 cm-l) is close to that for acetone in inert solvents such as cyclohexane (table l) so clearly there can be no significant hydrogen bonding in the cage. This result accords with conclusions from X-ray measurement^,^^ dielectric relaxation data 31 and proton n.m.r.When the concentration of ace-tone in liquid water is increased the absorbance of the band at 1697 cm-l falls and a new band at ca. 1708 cm-' grows in. This band is close to that observed for TABLE 1.-INFRARED DATA FOR ACETONE CARBONYL STRETCH IN VARIOUS SOLVENTS solvent band maximum/cm-' hexane 1721 carbon tetrachloride 1718 tetra hydro furan 1717 cyanomethane 1714 dimethyl sulphoxide 1709 methanol 1708 water (25 "C) 1697 water clathrate 1722 perfluoro-t-butyl alcohol 1688 acetone in methanol. It is easier experimentally to study the behaviour of acetone in low concentration as different aprotic cosolvents are added. This is the technique used in our work on R2N0 solutions and we refer to the acetone as a " probe " of the medium.When this is done with cosolvents such as cyanomethane we can again follow the loss of the 1697 cm-I band and growth of the 1708 cm" band. In solutions rich in MeCN this band loses intensity and a third band characteristic of acetone in pure MeCN grows in. Thus there are three species which we identify as the dihydrate the monohydrate and non-hydrated acetone. We conclude firmly that acetone sheds two hydrogen bonds on moving from liquid water into the clathrate. N.M.R. STUDIES As a check on the infrared results we have examined the 13C(CO)and 170n.m.r. spectra for acetone water systems. In our studies on triethylphosphineoxide (Et,PO) as a probe of mixed solvent systems we found that there was a good correlation between the 31Pchemical shifts and the P-0 stretch infrared shifts for a range of pure solvents which could be used to reconstruct the 31P shifts in mixed solvents.33 This was useful because the 31P spectra in mixed solvents are.rapid averages of the shifts for the different solvates involved whereas the infrared spectral features for all the sol- vates are in principle resolved. The 13C studies were very noisy despite long accumulation times but we were able to detect features for solutions in water at room temperature (the dihydrate) in partially frozen systems corresponding to the monohydrate and of course for pure acetone. Unfortunately the band for acetone in the solid clathrate was too broad to detect with our spectrometer. However all these bands were detected in the I7O spectrum including that for the clathrate and these correlate well with the infrared results (fig.3). The fact that the I7O resonance for the solid clathrate was no broader than the liquid-phase bands con- firms the remarkable rotational freedom of the encaged acetone molecules. M. C. R. SYMONS AND G. R. EATON r A0 -20 h Ob" lk E a W 3 x 0 -2 MeOHo -LO .3 0 /o frozen-30"~; 2 \:exone -60 hrate \ J I I I I 6 12 18 24 30 Av (carbonyl)/cm-' FIG.3.-Relationship between n.m.r. and infrared data for the carbonyl group of acetone in various states of solvation. ANCHORED US FREE CLATHRATES We need to ask why acetone changes its state so drastically on becoming enclath- rated.We stress that anchored guests are quite common and these include basic aprotic molecules such as amine~.~~ In our view the major reason for this change lies in the fact that in order that all the water molecules be tetrahedrally hydrogen- bonded in the clathrate cages the excess of (LP)freegroups must react with the (OH) groups bound to acetone Me,CO(HO) + 2(LP)fr,e +-Me2C0 + 2(0 * * HO) (2) where (0 * * HO) represents cage hydrogen bonds. Reaction (2) is favoured for acetone because it forms weak hydrogen bonds but evidently such a process is dis- favoured for most amines which form much stronger bonds. We conclude that acetone forms two weak hydrogen bonds in liquid water because of the availability of (OH)freegroups which are willing to form hydrogen bonds whenever possible.These are not available in normal clathrate cages where each water is tetrahedrally co- ordinated to four neighbours. INFRARED LINESHAPES The change in lineshape is in accord with these conclusions. The C=O stretch band in fluid aqueous solutions is Gaussian (fig. 2). This means that the width stems primarily from a range of subtly differing structures having long short and bent hydrogen bonds just as occurs for water-water interactions. In marked contrast the band for the clathrate is Lorentzian. This means that the width is kinetically controlled indicating considerable rotational freedom within these cage^.^^'^^ CAGE MODEL FOR ACETONE SOLUTIONS THE CASE FOR LIQUID-PHASE CAGES Our results show that if cages are important for liquid-phase systems they do not resemble those in the solid state and hence theories based on liquid-phase cages should not be justified on the basis that acetone forms a clathrate h~drate.~~~~~ What is important is that full allowance must be made for the fact that acetone forms two hydrogen bonds.In our view this reaction is likely to dominate most of the properties of aqueous acetone rather than cage formation. In the following we briefly consider alternative explanations for some observations that have been explained in terms of a tendency to form liquid-phase cages. EVIDENCE FROM RELAXATION STUDIES Both n.m.r.34*36 and dielectric 37 relaxation studies show that the correlation time (7,) for water molecules lengthens as basic aprotic cosolvents are added.This is interpreted in terms of increased water-water interactions and hence of cage formation. We wonder if the postulate of this structure-forming effect in the liquid phase is really necessary. We tentatively suggest that it is possible to understand the decrease in rotational freedom for water molecules in terms of eqn (2). We postulate that a major contribu- tion to rotational motion in normal water comes from the (OH)free water molecules. Their transformation into water bonded to acetone will inhibit such rotation. We stress that loss of (0H)free is balanced by a gain in (LP)f,,e units and hence we need to postulate that their rotational behaviour is less "free " than that of (OH)freeunits.This is reasonable since both protons remain hydrogen-bonded to bulk water in (LP)rreeunits. EVIDENCE FROM SPECTROSCOPIC STUDIES It has been suggested that the slight downfield shift 38 or plateau 25 that is initially observed in the water proton resonance as acetone is added is indicative of structure formation involving cages. However we have pointed out that this behaviour is a necessary consequence of the scavenging of groups.39 These groups must have their OH resonance close to that for monomeric water and hence they will con- tribute a considerable upfield shift to the time-averaged resonance. Their loss gives a downfield component that is almost balanced by other upfield shifts as we have dem~nstrated.~~ Similarly the loss of the band in the overtone infrared spectrum when basic aprotic solvents are added arises directly from the scavenging of groups by the base.It cannot be taken as proof of the formation of extra water structure. Again the ultraviolet shifts in the n+n* band for acetone in the water-rich region 40 can be fully explained in terms of the change from di-hydrogen-bonded carbonyl to mono-hydrogen-bonded and non-hydrogen-bonded units and in no sense requires the formation of cages. ADDITION OF ALCOHOLS Here we refer to some studies of the effect of added t-butyl alcohol and of methanol on the acetone clathrate. This was undertaken in the hope that with Me,COH unusual effects due to cage-cage affects might be detected. In fact nothing unusual was noticed.Methanol lowered the temperature at which the clathrates formed in the warming cycle [see experimental method (b)],whereas t-butyl alcohol raised it. M. C. R. SYMONS AND G. R. EATON 39 This presumably simply reflects the gain or loss of fluidity at low temperatures. The apparent decomposition temperature range was lowered for both systems and we were unable to form the clathrate at mole fractions (of alcohol) >0.016. BONDING IN T-BUTYL ALCOHOL CLATHRATES Finally we should mention a very interesting study of t-butyl alcohol clathrates by Gorbunov et aL41which in some ways resembles our own work. An infrared study in the fundamental 0-H stretch region for the mixed alcohol-H,S clathrate showed a narrow component in the 3608 cm-l region well removed from the absorption for the bound water molecules assigned to (OH)freeoscillators.It was reasonably concluded that the Me,COH molecules are not bonded (anchored) to the cage. This is not a necessary requirement since ROH molecules by forming two hydrogen bonds to water can knit into the cage structure with minimum disturbance. Indeed it is probable that this is the way ethanol molecules are en~lathrated.~~ On the other hand the structurally similar molecule Me,CNH, is freely enclathrated with no guest- host hydrogen These results which seem convincing suggest that Me,COH molecules like acetone undergo a dramatic loss of solvation on enclathration. 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