The study of solvated ion clusters provides detailed in-formation on microscopic interactions of ions with their sol-vent. Better ion–solvent and solvent–solvent interaction po-tentials may arise from experiments on small solvated–ionclusters. In studies of larger clusters, containing tens of sol-vent molecules, or more, the evolution of solvation from theconfined geometry to the bulk may be inferred.
Photoelectron spectroscopy~PES! is an important toolfor studying anion–solvent interactions, both in isolatedclusters1–5 and thin films of bulk solution.6–8 It measures thebinding energy of the anionic electron in the cluster, which isdirectly influenced by the ion–polar solvent interaction po-tential, distance, and orientation. Due to the vertical nature ofthe photodetachment, solvent–solvent interactions havehardly any effect on the energetics of the process.9 The pho-todetachment energetics can be calculated by molecular dy-namics simulations of ion solvation by switching off the in-teractions, which are related to the anion, or alternatively byremoving the highest occupied molecular orbital electron inab initio calculations. This feature enables direct comparisonof the PES binding energies with theory. Note that this fea-ture is unique to the studies of anion solvation, where theinitial state in the photodetachment process is theequilibrium-state solvated anion. There is no equivalent ex-perimental scheme for solvation studies of cations by photo-detachment.
In a previous work in our laboratory, Markovichet al.have studied the solvation of halogen anions in water clusters
by using PES,5 while Pereraet al. have used molecular dy-namics to simulate this system.10 On the basis of comparisonof the experimental results with different calculations, wededuced that each of the ions I2, Br2, and Cl2 is attached tothe surface of the water clusters.11 This behavior may berationalized in terms of balance between ion–solvent andsolvent–solvent interactions. In the case of anion–waterclusters, the solvent–solvent hydrogen bonding interactiondominates, and a maximal number of hydrogen bonds isformed, with only partial stabilization of the anion by thesolvent molecules.
Perera and Amar12,13 have analyzed the behavior of dif-ferent neutral solutes in argon clusters in terms of thestrength of solute–argon interaction relative to argon–argoninteraction. They found that relatively small solute–argoninteractions would produce a surface solvation state, while asthe solute–argon Lennard-Jones parameter2e exceeds 1.5times the Ar–Are parameter, the system forms an interiorsolvation state. Although systems of ions solvated by polarsolvents will probably not fall into the same category asthese weakly bound van der Waals complexes, the generalidea applies: As ion–solvent interaction becomes muchstronger than solvent–solvent interaction, the cluster willhave a larger probability of forming interior solvation states.This condition is fulfilled in anion–Arn clusters, where theanion-induced dipole interaction overcomes the week argon–argon interaction. PES experiments by Bowen andco-workers2 on O2Arn and by Neumark and co-workers onI2Arn ,
shown that these systems produce symmetric solvation shellstructures.
In this work we present photoelectron spectra of halides:Cl2, Br2, and I2, solvated in acetonitrile clusters—~CH3CN!n . Acetonitrile is commonly used as an aproticorganic solvent with good electrolyte solvation properties. Itwas chosen due to the combination of high dipole moment~3.92 D! and lower propensity to form hydrogen bonds.These properties result in favorable ion–solvent interaction~relative to water!, possibly leading to full solvation of theanions in the acetonitrile clusters, in contrast with the anion–water clusters’ behavior.
We analyze the results using classical molecular dynam-ics ~MD! simulations of halogen anions solvated in acetoni-trile clusters. A new potential function has been devised forthese simulations, since the existing potentials, used to de-scribe bulk acetonitrile properties, were not able to reproducethe experimental ion–solvent electrostatic interactions prop-erly. Berkowitz and co-workers10 have shown in their workon the anion–water clusters, that polarization effects in suchsystems play a very important role in the energetics andstructure of the clusters. We have therefore modified the sixsite model for acetonitrile by Bohm, McDonald, andMadden15 to include polarization.
II. METHODS
A. PES experiment
The experimental setup is described in detail elsewhere.5
The clusters are formed in a pulsed supersonic expansion of1–2 bar of Ar, 40 mbar of CH3CN, and about 1 mbar of ahalogen-containing compound~CCl4, C6H5Br, or C7H7I!.Between 100 and 200 eV electrons are injected into the earlystages of the expansion in order to form the charged clusters.The clusters are cooled by further flow in the expansion, passthrough a skimmer to a second vacuum chamber, and aremass separated by a reflectron time-of-flight~TOF! massspectrometer.
Prior to photodetachment the ions are mass selected byan electrostatic mass gate, and an electric field impulse isused to slow the ions down to;20–100 eV kinetic energy inorder to reduce Doppler broadening. About 2 cm down-stream, a vacuum ultraviolet laser pulse intersects the ionbeam, and the kinetic energy distribution of the detachedelectrons is analyzed in a 250 cm ‘‘magnetic bottle’’ TOFphotoelectron spectrometer.16
All the spectra are taken with 7.9 eV photons, generatedby F2 excimer laser. The spectrometer is calibrated withspectra of the halogen anions~Cl2,Br2,I2!.17
B. MD simulation
We have tried to use the six-site model of Bohm, Mc-Donald, and Madden,15 which contains a partial charge and aLennard-Jones center for each of the atoms in the acetonitrilemolecule but does not include polarization terms. This modelis able to account for liquid acetonitrile properties, such asneutron scattering results. Del-Mistro and Stace18 made use
of these potential parameters in simulations of the IR fre-quency shifts in small acetonitrile clusters as measured byBuck and Ettischer.19 Their calculations maintain reasonableagreement with the experiment and reproduce the samestructures which were deduced from the experimental workof Schermann and co-workers.20 We could not fit our PESexperimental results with this model, apparently due to theexistence of strong polarization effects in our charged clus-ters.
In our calculations we have kept the Lennard-Jones pa-rameters of the Bohm potential, yet modified this potentialby making the following two changes.
~i! The introduction of a new charge distribution whichrests on a more recentab initio calculation onacetonitrile.21 In this work, except for the acetonitrileitself, the association energy of Cl2~CH3CN!1 wascalculated. It was in good agreement with the experi-mental result by Hiraoka, Mizuse, and Yamabe.22 Thecharge distribution from this calculation overesti-mated the acetonitrile dipole moment~4.42 D!. Wehave uniformly reduced the charges by 10% toroughly match the experimental value of the gas-phase dipole moment of acetonitrile~3.98 D versusthe experimental 3.92 D!. This procedure is similar tothe construction of the POL model for water,23 wherethe charge distribution was also reduced by a constantfactor upon the inclusion of polarization effects in thepotential.
~ii ! The introduction of polarization effects by attributingpolarizability to the anion and to each atomic site.
In our model, anion potentials also include Coulomb,Lennard-Jones, and polarizability terms. The parameters forthe anions were taken from halide–water simulations. Someof the Lennard-Jones radii were slightly corrected in order toobtain a better fit to the experiment. The Lennard-Jones pa-rameters of the anions~especially the diameter! constitute themain difference between the different halides, within themodel that we use. The results of the simulations and their fitto the experimental values are, therefore, very sensitive tothe magnitudes of these parameters.
The simulation method was similar to the one used pre-viously for the halide–water system.10 The polarization en-ergy could be calculated both in a self-consistent manner asin the POL model for water23 ~iterative calculation!, or ap-proximately~one-step calculation!, by using the electric fieldgenerated at a site by the permanent atomic charges of theother molecules~or ion! to calculate the polarization energyat the site, neglecting the correction to this field by the otherinduced dipoles. We noticed, that the values of the polariza-tion energy calculated in the two methods differed by a fewpercent. This difference is negligible relative to the total en-ergy of a cluster. Therefore, we have decided to use theeconomical method, using the approximate values of theelectric field.
We did not include a three-body interaction term~solvent–solvent–ion! since, as in the case of the halide–water simulations, its contribution to the total energy of a
2676 Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
cluster is probably small, and the parameters of such a termfor acetonitrile were not available to us.
The total potential energy of a cluster is given by
U tot5Uel1ULJ1Upol , ~1!
where
Uel5(i. j
qiqjr i j
, ~2!
ULJ5(i. j
SAi j
r i j122
Ci j
r i j6 D , ~3!
Upol521
2 (i
a iEi2, ~4!
where the Coulomb field,Ei , at each site is given by:
Ei5(j
qjr i j2 r̂ i j . ~5!
The sums are over all the sites, or site pairs of differentmolecules, in the cluster.Ai j andCi j are either homoatomicor heteroatomic Lennard-Jones coefficients,r ij are inter-atomic vectors,qi are the site charges, anda i—site polariz-abilities.
The parameters used in the simulations are listed inTable I.
In the simulations we calculate the stabilization energyof the anion by the solvent,Estab~MD!, which is the quantityto be compared with experimental results. It is evaluated asthe difference between the initial-ionic state and the final-neutral state~at the same nuclear coordinates! which is theresult of a vertical electron detachment transition:
Estab~MD)5~Ueli-s1ULJ
i-s1Upion!2~ULJ
n-s1Upntl!, ~6!
where the superscript i-s means ion–solvent potential energyand n-s means neutral halogen–solvent potential energy,Upion is the total polarization energy of the ionic cluster, and
Upntl is the total polarization energy of the neutral cluster.The experimental temperature, which is cluster-size de-
pendent, is estimated roughly by using the evaporative en-semble model24 following Lisy and co-workers.25 The inputsto this calculation are stepwise binding enthalpies, measuredexperimentally for the small clusters by Hiraokaet al.,22 andthe evaporation TOF between the cluster source and photo-
electron spectrometer. For I2~CH3CN!n clusters, the esti-mated upper bound for the temperature ranges from about160 K for n54 to about 70 K forn550.
In order to explore as many configurations as possiblewithin a limited computational effort, we perform severalheating–cooling cycles during a MD run. The initial configu-rations are equilibrated at 200–300 K for a period of 50–1000 ps~depending on cluster size!. Subsequently, clustersare cooled to the estimated experimental temperature, run forabout 400 ps, and periodically heated to 200–300 K for briefperiods of about 40 ps. Since both the configurations of theclusters andEstab are temperature dependent, data are col-lected only during the colder parts of the trajectories.
III. RESULTS AND DISCUSSION
A. Experimental observations
We have measured the PES of halogen anions solvatedin n acetonitrile molecules, wheren51–33, 1–40, and 1–55for Cl2, Br2, and I2, respectively. The spectra of I2, Br2,and Cl2 clustered with~CH3CN!n are shown in Figs. 1, 2,and 3.
In our analysis we postulate that the peaks of these spec-tra represent the average vertical binding energies of elec-trons in the clusters@BEn(n)#. We also assume that in theseclusters the excess charge is practically localized on the an-ions. This hypothesis is supported by the recent experimentsof Johnson and co-workers.26 We estimate the uncertaintiesin the vertical binding energies to be of the order of 15 meV~for small clusters withn51–3! up to 150 meV@for thelargest X2~CH3CN!n clusters#. The difference between thevertical binding energy of a cluster and the electron affinityof the bare ion is defined asEstab(n):
Estab~n!5BEn~n!2BEn~0!. ~7!
Estab(n) is essentially the electrostatic stabilization of thesolvated anion in its equilibrium configuration. In the verticalphotodetachment process, the neutral halogen–solvent dis-tances in the final state are identical to these of the anion–solvent, making the neutral atom–solvent interaction rela-tively weak. The simulations give neutral–solvent interactionenergies which are up to 4% of the ion–solvent stabilizationenergy. Our experimental values ofEstab(n) are listed inTable II together with the values ofEstab~MD!.
The scaled experimentalEstab(n) values for the threehalides are plotted in Fig. 4 as a function of cluster size. The
TABLE I. Potential parameters used in the simulations.
HC
methylC
nitrile N Cl2 Cl Br2 Br I2 I
q (e) 0.220 20.582 0.332 20.410 21 0 21 0 21 0s i i
a ~Å! 2.1 3.0 3.4 3.3 4.32 3.7 4.7 4.0 5.15 4.63e i i
aThe transformation from these atomic parameters toAi j andCi j , used in Eq.~3! aree i j5Ae i ie j j ; s i j5(s i i1s j j )/2; Ai j54e i js i j12; Ci j54e i js i j
6 .
2677Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
scaling was done by dividingEstab(n) for each halide byEstab~1! of the same halide. In this way, a unified scale, inunits of one acetonitrile–anion stabilization energy, is ob-tained. Note that iodide and bromide data fall on the samecurve~within the experimental error!, while the chloride sta-bilization energy steps are diminishing faster. The deviationof the chloride curve from the other two is mainly due to theabnormally highEstab~1! of the chloride. The PES stabiliza-tion energy of the complex of Cl2~solvent!1, both in waterand acetonitrile, is inexplicably much higher than the mea-sured binding enthalpy of the clusters.22 The binding en-thalpy for such a small cluster should be close to the stabi-lization energy due to the absence of solvent–solventinteractions.27 Also note that in the case of the Cl2~CH3CN!1cluster, we were unable to reproduce the PES stabilizationenergy in the MD simulations~0.24 eV deviation!, by usingthe same potential parameters which fitted well~within 0.05eV! the rest of the clusters. We suggest that in this clusterthere is a very strong spectroscopic effect in the PES, result-ing in intramolecular vibrational excitation in the neutral
cluster. Such an effect would not be reproduced in our simu-lations which use rigid solvent molecules. Therefore, wetried to scale the chloride curve by dividing theEstab(n)values by half of theEstab~2!. In this case, Fig. 4 reveals agood fit between the curves of the three halide–solvent clus-ters up ton59.
A general trend in the scaledEstab(n) of all clusters isthe gradual decrease in the differential stabilization energy,DEstab(n), upon the stepwise addition of solvent molecules,DEstab(n)5Estab(n)2Estab(n21). This trend is well illus-trated in Fig. 5. We attribute the decrease in differential sta-bilization to the influence of solvent–solvent interactions onthe structure, which results in reduced ion–solvent interac-tion energy.
A common behavior to the three curves in Fig. 4 is theleveling of Estab(n) beyondn512 for I2, n510–11 forBr2, andn59–11 for Cl2. This abrupt change inEstab(n)was not observed with other solvents~H2O, CO2, Ar! inother PES experiments. We associate this phenomena withthe completion of the first solvation layer around the halide.The absence of further stabilization of the anion with clustersize, reflects the switching in the relative strength of anion–solvent versus solvent–solvent interactions. In the first sol-vation layer, the relatively strong ion–solvent interactions,
FIG. 1. Photoelectron spectra of I2~CH3CN!n in the range ofn51–55,taken with 7.9 eV photon energy.
FIG. 2. Photoelectron spectra of Br2~CH3CN!n in the range ofn51–40,taken with 7.9 eV photon energy.
2678 Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
which govern the cluster structure, result in efficient anionstabilization. In contrast, in the second layer, solvent–solventinteractions defeat the much weaker ion–solvent interac-tions, contributing no further stabilization to the anion.
At n59–20, the chloride curve is somewhat lower inscaled stabilization energy than the other two. This is prob-ably due to the smaller number of solvent molecules in thefirst layer around the smaller chloride.
Changes in the character of the clusters are reflected alsoin the mass spectra of the X2 ~CH3CN!n . There is a sharpdrop in cluster abundance beyondn512, 9, and 9 for I2,Br2, and Cl2, respectively~see Fig. 6!. Although the shapeof cluster size distribution may be changed according to clus-tering conditions at the source, these sharp drops do not de-pend on source conditions. These sharp changes in abun-dance in the mass spectrum will be discussed in relation toour MD simulations.
B. Simulation results
The MD simulations provide a rationalization for thedecrease in differential stabilization energy with increasingcluster size in the small clusters regime~Fig. 5!.
Maximum stabilization of the halogen anion with oneacetonitrile is fulfilled in a geometry in which the symmetryaxis of acetonitrile is normal to the anion sphere with allthree hydrogens equally distant from the anionic center.28
The clusters withn52–4 maximize ion–solvent interaction
by having the same ion–solvent configurations as inn51while minimizing repulsion between the ligands. The result-ing structures are antiparallel, collinear forn52, planar–triangular structure forn53, and roughly tetrahedral struc-ture for n54. In Fig. 7 we display snapshots from thesimulations of a few small I2~CH3CN!n clusters.
For larger clusters (n.4), as the congestion of ligandsaround the anion increases, the acetonitrile ligands bendfrom the normal to the anion surface and the solvent–solventinteraction becomes attractive, with the penalty of a reduceddifferentialEstab. For instance, the solvent–solvent Coulombenergy for I2~CH3CN!4 is1160 meV, while for I2~CH3CN!6it is already280 meV. This change can be noted in theaverage ion–hydrogen radial distribution function for then54,5 clusters in Fig. 8. The behavior is common to thethree halides: Atn54 all three hydrogen atoms of each ac-etonitrile molecule are nearly equivalent. Atn55 two hy-drogen atoms become closer to the anion while the third isfurther away, interacting with the nearest-neighbor nitrogenatom, and two separate peaks appear in the ion–hydrogendistance distribution function. In larger clusters, some of themolecules have one hydrogen in contact with the anion andthe other two pointing toward the second solvation layer.
The I2~CH3CN!6 and Br2~CH3CN!6 clusters possess acommon structure: The acetonitrile molecules form a planarring around the ion, having head to tail interactions betweenthe molecules, and one or two hydrogen atoms in contactwith the ion, optimizing both solvent–solvent and ion–solvent interactions. Cl2~CH3CN!6 is too crowded to form aring and the molecules move out of the plane of the ring.
Extreme care has to be taken in the simulations of largerclusters. Depending on the way the trajectories are run, dif-ferent structures may result. The way the clusters are an-nealed in order to obtain the ‘‘experimental’’ structure iscrucial to both structure andEstab. One important example inthis work is the I2~CH3CN!12 cluster. On running a longtrajectory~a few ns! at 110–120 K, which is the estimatedupper bound to the experimental temperature, with periodicheating to 200 K, the cluster would have one solvent mol-ecule at the second solvation shell, on the average. No spe-cial order appears in the first solvation layer and the experi-mental stabilization energy is reproduced within 40 meV. Ifone uses very short~2–3 ps! heating periods to 300 K, whichwould cause dissociation on a longer time scale, andquenches them to a relatively low temperature~70 K!, then amuch more stable structure appears. This ‘‘crystalline’’structure is about 200 meV lower in total energy than theformer ‘‘amorphous’’ structure, while havingEstab~12!which is smaller by about 200 meV than the experimentalone. It has all the 12 molecules in the first layer, formingdistorted cyclic trimers which are arranged in a symmetricshell, having aT point group symmetry~Fig. 9!. This stablestructure can survive heating up to 150 K for several nano-seconds.
Another example isn511. Regular trajectories with pe-riodic heating produce equilibrium configurations having onesolvent molecule in the second layer of I2~CH3CN!11. Whenan initial configuration of this cluster is produced by remov-
FIG. 3. Photoelectron spectra of Cl2~CH3CN!n in the range ofn51–33,taken with 7.9 eV photon energy.
2679Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
ing one molecule from the more stable form of theI2~CH3CN!12, or by quenching from very high to low tem-perature~70 K!, a more stable form of the cluster is obtained.This isomer has all the 11 solvent molecules in the first sol-vation layer. It is lower by about 100 meV in total energythan the one obtained in the regular runs, and again, whilethe stabilization energy obtained from the regular run fits theexperimentalEstab~11! within 30 meV, this quantity issmaller in the more stable isomer by 120 meV.
We have also tried to find a more stable configuration ofthe I2~CH3CN!13 cluster by taking the most stablen512cluster, adding a molecule to the second solvation layer, andrunning for a few nanoseconds at 110 K. The additional mol-ecule moved into the first layer, producing the same disor-dered structure that was achieved previously in the regularheating–cooling cycles. Note, that on going from the stableform of the I2~CH3CN!12 cluster to the I
2~CH3CN!13 cluster,the total energy of the cluster is lowered by 320 meV, whileon going fromn511 to 12 it lowers by 450 meV. Thechanges in the solvent–solvent interaction energy are moreilluminating: It becomes more negative by about 400 meV
between the most stable forms ofn511 and 12 while nochange is observed betweenn512 and 13.
Looking for a similar behavior in the bromide and chlo-ride clusters, we found that both Br2~CH3CN!9 andCl2~CH3CN!9 clusters, as in the case of the I2~CH3CN!12,consist of distorted cyclic trimers of acetonitrile. In bromideand chloride withn59 the acetonitrile cage is not completeand the symmetry is lower. A snapshot of the Cl2~CH3CN!9cluster is shown in Fig. 10. In this case, the changes in totalcluster energy betweenn58 and 9 and betweenn59 and 10are almost constant, while the changes in the solvent–solventenergy are those which contain valuable information. Both inbromide and chloride, on going fromn57 to 8 and fromn58 to 9, the solvent–solvent energy becomes 300–400meV more negative, while on going fromn59 to 10, it isreduced by less than 100 meV. We believe that the intensitydrops in the mass spectrum of the X2~CH3CN!n clusters arerelated to these observations. The ordered cyclic trimer struc-tures of the solvent in the Cl2, Br2, I2 clusters withn59, 9,and 12, respectively, have relatively low total energy andvery low solvent–solvent energy which makes the enthalpy
TABLE II. Experimental and simulatedEstab values of the X2~CH3CN!n clusters. The temperatures at which the simulatedEstab values were collected, andthe number of solvent molecules in the first solvation layer are indicated. We estimate the errors in determining experimentalEstabto be of the order of 15 meVfor n51 up to 150 meV forn555. Uncertainties in simulatedEstabare of the same order of magnitude. The numbers in parenthesis refer to stable structureswhich are hardly accessible through MD trajectories.
of association of these clusters more negative with respect tothe other, amorphous, clusters. Consequently, these clustersare less susceptible to evaporation of acetonitrile molecules,leading to a ‘‘magic number’’ in the mass spectrum. Aspointed out recently by Sremaniak, Perera, and Berkewitz29
structures and phase transitions in such systems are often
governed by solvent–solvent interactions, while the proper-ties of the ion play a secondary part in such phenomena.
In discussing the reliability of our simulations, one mustbear in mind that these structural properties are not ex-tremely sensitive to the exact form of the potential used inthe simulation. Increasing the atomic charge distribution by5%, or taking in or out the polarization term, did not changethe structures significantly.Estab(n) values, however, werevery sensitive to the potential parameters that were used anddid not easily fit the experimental ones. Furthermore,Estab(n) was found to be very sensitive to the exact configu-ration of the molecules in the first solvation layer, which maychange upon changing the temperature of a cluster or trap-ping in different local minima. Further discussion of phasetransitions and temperature effects in MD simulations ofhalide–water clusters can be found in Ref. 29.
Beyond the first solvation layer (n.12), theadditionalacetonitrile ligands form disordered structures, maximizingthe interaction with the molecules in the first layer andamong themselves. These molecules usually form nonlinearhead-to-tail configurations and less frequently—antiparallelpairs. The dipole moments of the molecules in the secondlayer are roughly perpendicular to the ion–molecule vector,adding no stabilization to the anion. They may rather distortthe structure of the first solvation layer, yielding negativestabilization of the anion. This is a fine example of ‘‘struc-ture breaking’’ geometries which is traditionally attributed tosecond solvation layers in the bulk.
FIG. 4. Scaled stabilization energies as a function of cluster size forX2~CH3CN!n clusters where X5Cl, Br, and I. The scaling was done bydividing Estab(n) by Estab~1! for each halide. For chloride, a scaling bydivision by Estab~2!/2 is also presented~full circles!—see the text for anexplanation.
FIG. 5. Scaled differential stabilization energies as a function of cluster sizefor X2~CH3CN!n clusters where X5Cl, Br, and I. The scaling was done bydividing @Estab(n)2Estab(n21)# by Estab~1! for each halide.
FIG. 6. Typical mass spectra of X2~CH3CN!n cluster series where X5Cl,Br, and I. Note the sharp drops in intensity beyondn59, 9, and 12 for Cl2,Br2, and I2, respectively. These drops are unique throughout the series anddo not disappear upon changes in cluster size distribution.
2681Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
The largest clusters we have simulated, I2 and Br2 with25 solvent molecules, and Cl2 with 23 molecules, had 12–13, 12, and 11 solvent molecules, respectively, in the firstsolvation layer at 90–110 K, having the outer solvent mol-ecules concentrated over a limited solid angle around the
anion. A typical snapshot of the I2~CH3CN!25 cluster is pre-sented in Fig. 11. This trend may probably continue in largerclusters: The second solvation layer might not be completed,building an acetonitrile drop which is attached to the firstsolvation layer. The anion–solvent stabilization added by themolecules beyond the first solvation layer is small relative tosolvent–solvent interaction energy. Therefore, the solventwould rather form a drop and not fill the second layer. Thesmall gradual increase in experimentalEstab(n) beyondn520–25 can be attributed to the electrostatic polarizationof the solvent drop beyond the first solvent shell~Born en-ergy!.
The evolution ofEstab(n) in halide–acetonitrile cluster isdistinctly different from that of halide–water clusters~Fig.12!. Note, that the water clusters do not show any leveling ofEstab(n) ~only a change in slope atn56!. Already at smallsizes, the halide–water clusters have structures that resemblepure water structures,5,29 while the halide–acetonitrile clus-ters probably need tens of solvent molecules to start buildingsolvent–solvent structures which resemble the acetonitrilebulk structure. In the X2~CH3CN!n clusters, the first solva-tion layer of the clusters is full~for n.12!, with 11–13
FIG. 7. Structures of small I2~CH3CN!n clusters, calculated in the MD simulations. The estimated experimental temperatures, at which the data werecollected, are indicated. van der Waals radii of atoms in acetonitrile are scaled down by 0.3 for visual clarity.
FIG. 8. I2–H and I2–C~methyl! radial distribution functions, accumulatedover 0.5 ns runs of I2~CH3CN!n clusters withn54,5 atT5160–150 K.
2682 Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
solvent molecules in this layer, while in halide–water clus-ters, the first solvation layer is only partially filled, with amuch smaller solvation number.
The marked difference in the structure of the solvent inthe first solvation layer, between water and acetonitrile, fadesaway on moving beyond the first layer. There, as in any polarsolvent, the solvent–solvent interaction becomes larger thanion–solvent interaction and the system would maximizesolvent–solvent interactions to form a densely packed dropof the solvent, rather than a spherical symmetric solvationshell around the ion. Thus, large ion–acetonitrile clusterswill have a solvent drop attached to a full first solvationlayer, while the water clusters may have the anion exposed atthe surface of the drop. At large enough clusters and high
enough temperatures, the entropy and long range polariza-tion energy may finally drive the anion toward the center ofthe cluster.
IV. CONCLUSIONS
We have demonstrated the sensitivity of PES to the mi-croscopic details of ion solvation. The pure ion–solvent elec-trostatic interaction, extracted from the PES experiment,shows significant differences between water and acetonitrile,reflecting differences in cluster solvation structures.
Classical molecular dynamics simulations, if used care-fully, are indispensable in the analysis of results, in spite of
FIG. 9. The structure of the highly stable form of the I2~CH3CN!12 cluster.Hydrogens are not shown in the picture for visual clarity. Nitrogen is rep-resented by the clear circles. All the acetonitrile molecules are combined incyclic trimers. Two viewing angles are presented: On the top, one of thefour C3 axes of theT point group is perpendicular to the plane of the paper.On the bottom, one of the threeC2 axes.
FIG. 10. The structure of the Cl2~CH3CN!9 cluster. As the I2~CH3CN!12cluster, it is built out of cyclic acetonitrile trimers, but this time they do notform a full solvation shell. Hydrogens are not shown in the picture for visualclarity. Nitrogen is represented by the clear circles.
2683Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
the rough potential functions employed. By finding anatomic charge distribution that produces better agreement be-tween the simulation results and the experiments, we mighthave improved the potentials used so far in acetonitrile simu-lations. The anion–acetonitrileEstabcalculations pose a morestringent test on the atomic charge distribution, in compari-son with the pure acetonitrile simulations.
As revealed to us in the case of I2~CH3CN!12, findingthe global minima of multidimensional systems such as ion–acetonitrile clusters with tens of solvent molecules, usingMD, is a difficult, and sometimes unreliable task. It is pos-sible, that the exploration of phase space toward the globalminimum of such a complex system may be done more ef-ficiently using other methods, employing random configura-tion changes in the search scheme. MD may be successfullyapplied afterwards, using the global minima as a startingpoint, to study dynamics, temperature effects, and structuresof complex systems.
ACKNOWLEDGMENTS
The research was supported by the James FranckGerman–Israeli Binational Program in Laser Matter Interac-tion, and by the US–Israel Binational Science Foundation.G.M. is a Clore Scholar. The help of I. Schek in running thesimulations is highly appreciated.
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FIG. 11. A typical snapshot of the I2~CH3CN!25 cluster atT590 K. Notethat the molecules outside the first solvation layer are concentrated to forma drop that is attached to the first layer. Hydrogens are not shown in thepicture for visual clarity. Nitrogen is represented by the clear circles.
FIG. 12. Scaled stabilization energies of I2~CH3CN!n and I2~H2O!n clusters
as a function of cluster size. For each solvent, theEstab(n) values are nor-malized toEstab~1!.
2684 Markovich et al.: Solvation of Cl2, Br2, and I2 in acetonitrile
