Sink or Swim: Ions and Organics at the IceâAir Interface - Journal of

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Sink or Swim: Ions and Organics at the Ice−Air Interface Arpa Hudait, Michael T. Allen, and Valeria Molinero* Department of Chemistry, The University of Utah, 315 South 1400 East, Salt Lake City, Utah 84112-0850, United States S Supporting Information *

ABSTRACT: The ice−air interface is an important locus of environmental chemical reactions. The structure and dynamics of the ice surface impact the uptake of trace gases and kinetics of reactions in the atmosphere and snowpack. At tropospheric temperatures, the ice surface is partially premelted. Experiments indicate that ions increase the liquidity of the ice surface but hydrophilic organics do not. However, it is not yet known the extent of the perturbation solutes induce at the ice surface and what is the role of the disordered liquid-like layer in modulating the interaction between solutes and their mobility and aggregation at the ice surface. Here we use large-scale molecular simulations to investigate the effect of ions and glyoxal, one of the most abundant oxygenated volatile organic compounds in the atmosphere, on the structure, dynamics, and solvation properties of the ice surface. We find that the premelted surface of ice has unique solvation properties, different from those of liquid water. The increase in surface liquidity resulting from the hydration of ions leads to a water-mediated attraction of ions at the ice surface. Glyoxal molecules, on the other hand, perturb only slightly the surface of ice and do not experience water-driven attraction. They nonetheless accumulate as dry agglomerates at the ice surface, driven by direct interactions between the organic molecules. The enhanced attraction and clustering of ions and organics at the ice surface may play a significant role in modulating the mechanism and rate of heterogeneous chemical reactions occurring at the surface of atmospheric ice particles.

1. INTRODUCTION Ice naturally presents a disordered, premelted liquid layer at the ice−air interface.1−7 This disordered layer is the locus for chemical reactions8−10 and plays a role in determining the rate of growth of snow crystals11−13 and the transport and electrification of charged species during the collision of ice particles.1,14,15 The thickness of the premelted layer diverges logarithmically on approaching water’s melting point.7,16,17 At temperatures relevant to ice clouds, 150−260 K, the premelted layer does not fully cover the ice surface, which results in an average thickness that is less than the width of one water molecule.7,16,17 Experiments indicate that solutes at the surface of ice can perturb the premelted layer to different degrees, depending on their hydrophilicity.18−23 While hydrophilic volatile organics (e.g., acetone and acetic acid) do not affect the premelted layer,19,22 ions increase the amount of premelted water on ice.20,24 To our knowledge, it has not yet been quantified how the liquidity, defined as the number density of liquid-like molecules at the ice surface, changes with the concentration of ions or hydrophilic organics, and what is the effect of the liquidity of the premelted layer on the mobility and interactions between solutes at the ice surface. Addressing these questions is the goal of the present study. Ions at the ice−vapor interface are squeezed between two solutophobic phases, because vapor cannot stabilize the ions and solvation of ions within the ice crystal is also unfavorable.25,26 The solubility of ions and molecular solutes in ice crystals is negligible: solutes are expelled from water as it crystallizes, accumulating at ice interfaces, either in grain boundaries or at the ice−air interface.18,23,27 Adsorbed ions at © 2017 American Chemical Society

the ice−vapor interface can be solvated by the already present disordered premelted water, or they can induce ice melting to create liquid-like water in their solvation shells. This suggests that the solvation of ions at the ice−vapor interface could be controlled by a balance between free energy of ion solvation, free energy of melting of ice, and interfacial free energy between liquid, ice, and vapor phases. A simulation study of ions at the liquid−vapor interface shows that ions experience stronger attraction at the interface than in bulk liquid water.28 The attraction between ions at the liquid−vapor interface is driven by the reduction of capillary deformations of the liquid surface as the ions are driven out of the interfacial region.28 A pair of ions experience attractive forces at the liquid−vapor interface, irrespective of whether they have the same or opposite sign charges or the solutes have no charge at all.28 While attraction between ions driven by reduction of capillary deformations may be also expected for ions pulled out of the ice−vapor interface, it is not known whether water-driven interactions exist between ions and other solutes naturally adsorbed at the ice−vapor surface. Oxygenated organic molecules, such as acetone, glyoxal, and acetic acid, are present in the atmosphere in trace amounts and adsorb on environmental ice surfaces.21,29,30 Experiments indicate that oxygenated organics do not increase the liquidity of the ice surface at temperatures ranging from 218 to 240 K.19,22 To our knowledge, experimental results are not available at higher temperatures. Different from the ions, oxygenated Received: May 23, 2017 Published: June 29, 2017 10095

DOI: 10.1021/jacs.7b05233 J. Am. Chem. Soc. 2017, 139, 10095−10103

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Journal of the American Chemical Society

their interaction with ice is too weak to result in significant adsorption to the surface at 260 K.74 2.2. Simulations. The simulations are evolved with molecular dynamics simulations using LAMMPS.75 The equations of motion are integrated with the velocity Verlet algorithm using a time step of 10 fs, except for the simulations involving glyoxal where a 5 fs time step was used. The temperature and pressure are controlled with a Nose− Hoover thermostat and barostat with time scales of 25 and 50 ps, respectively. To investigate the behavior of solutes at the ice surface we construct a simulation cell that exposes the basal (0001) plane of hexagonal ice to the vapor phase. The simulation cell is 14 water layers deep, with the two bottom layers fixed, and exposes a surface (xy plane) of 10 nm × 10 nm to the vapor. The simulation cell is nonperiodic in the z direction; a reflective wall at about 2 nm from the ice surface terminates the gas region. The simulations were evolved in the NpT ensemble with pressure 1 bar controlled independently in the two directions tangential to the ice−vapor interface (xy plane). We do not apply a barostat in the z-direction, as it is open to the gas phase. To create a system with solutes (ions or glyoxal) at the ice−vapor interface, we first randomly place the solutes 3 Å above the ice surface and equilibrate the systems for 100 ns. The equilibrated configurations are then further evolved for 1 μs. The largest cluster of ions at the ice surface is identified with a cutoff distance of 0.75 nm. The distribution of the largest cluster sizes is computed over 900 ns for 35 ions at 260 K. We identify liquid molecules in these microsecond long simulations with the CHILL+ algorithm.76 We count the liquid molecules in the 12 Å of ice close to the vapor interface, which is the maximum depth of the perturbation induced by the ions at the ice surface. 2.3. Calculation of Properties. To compute the potential of mean force at the ice−vapor interface, we place 2 ions or two glyoxal molecules at 3 Å from the ice surface, and we first equilibrate the system for 5 ns before collecting statistics from 3 μs simulations. The radial distribution function, g(r), between two solutes is computed from the histogram of the distances between the two solutes in the 3 μs simulation. The distance between the glyoxal molecules is calculated as the distance between the centers of the carbon−carbon bond in each molecule. Since the solutes can only graze the surface of ice, we consider that they move in a cylindrical shell of height 7.37 Å (2 ice layers), the typical depth a solute penetrates from the ice−air interface. The free energy of interaction between the solutes at the surface relative to the free energy of the infinitely separate pair, ΔGsurface(r) is then computed in the standard way from the radial distribution function, ΔGsurface(r) = −RT ln g(r). We characterize the spatial extent of the shell of disordered water around the ions through the calculation of the averaged density profile of disordered water around a pair of ions separated at various distances. To compute the profiles, we perform simulations in which a pair of ions is placed at the ice surface at 260 K, restrained at fixed distances ranging from 5 to 40 Å through a harmonic spring with force constant 4 kcal mol−1 Å−1. We evolve each of these simulations over 400 ns, of which the last 300 are collected for the calculation of the local density of liquid water across the ice surface. To compute the latter, we grid the surface in 0.5 × 0.5 Å2 bins and average the number of water molecules classified as liquid in each bin. We estimate the length scale of interference of the solvation shells of the ions, 1 nm, from the time averaged density profile of liquid-like molecules (see section 3.2). We compute the difference in the number of liquid-like molecules at a given ion separation compared to the two ions at infinite distance, ΔN(r) = N(r) − N(∞), from the simulations of constrained ion pair by measuring the number of molecules within the length scale of interference of the solvation shell of the ion, 1 nm, for ion separations ranging from 5 to 40 Å, using 40 Å as a proxy for infinite distances. ΔN(r) is used to compute ΔGmelt(r), which is extrapolated to the locus of the free energy minimum, 4.7 Å using a second-order polynomial with the expression, ΔGmelt(r) = −1.4659 + 0.1279r − 0.0026r2 to fit the computed ΔGmelt(r) in the range of 5−17 Å. Based on the results discussed in section 3.2 below, we approximate the liquid region generated by an individual ion as a disk of radius 1 nm. When the ions are farther than 2 nm apart, the pair of ions at the

organics are volatile. A number of experimental and theoretical studies indicate strong enthalpy of adsorption ranging from −70 and −30 kJ mol−1 for oxygenated organic compounds at the ice surface,31−41 which support strong adsorption of organics at the ice−air interface. However, the hydration enthalpies of organic molecules are a small fraction of the ones for ions, e.g., the hydration enthalpy of acetone is less than 3% of the ones for NaCl or LiCl.42,43 Experiments and simulations indicate that adsorbed polyaromatic hydrocarbons aggregate at the ice surface.44,45 It has not yet been elucidated how the surface of ice modulates the interaction between hydrophilic oxygenated organics at the surface of ice. Experiments revealed an increase in the adsorption of hydrophilic organics at the ice− vapor interface in the presence of the strong acid HCl.24 It is an open question whether the surface disorder induced by ions impacts the aggregation of hydrophilic organics at the surface of ice. Here we use large-scale molecular dynamics simulations to characterize the effect of ions and glyoxal, one of the most abundant oxygenated organics in the atmosphere,46−49 on the liquidity of the ice surface, and the role that liquidity has on the mobility of solutes and their free energy of association at the ice−air interface at temperatures relevant to the troposphere. Our results indicate that hydrated ions cluster at the ice surface driven by a long-ranged water-mediated attraction that originates from the drive to minimize the liquidity created by the ions at the surface of ice. We find that glyoxal also aggregates at the ice surface, but the clusters are dry and driven by direct interactions between the solute molecules. The different effect of ions and organics on the liquidity of the ice surface results in opposite effects on their surface mobilities. This study indicates that the ice surface influences the solvation, aggregation, and mobility of ions and organics in a way that is distinct to that in liquid water.

2. METHODS 2.1. Models. We model water using the monatomic model mW,50 which accurately reproduces the thermodynamics,51−60 structure,50−53,61−65 and premelted layer5,16,66 of ice at less than 1% of the computational cost of atomistic models.50,57 The ions are represented by solute S,67 a strongly hydrophilic monatomic solute that has strong attraction with water. The S ions interact among themselves and with water through short-range potentials, yet they reproduce the effect of LiCl (a pair of S ions is equivalent to one LiCl) on the structure of liquid water, melting temperature of ice, crossover between crystallization and vitrification as a function of salt concentration, and the activity coefficient of water in solutions.62,67−69 The S-mW model reproduces well (section 3.2) the free energy of attraction between ions in atomistic simulations of dilute LiCl solutions, which predict an almost inaccessible contact pair (CP) and weakly stabilized solvent-separated pair (SSP) at 0.47 nm.70 We also test the effect of less strongly interacting ions Sweak, which have the same size and ion−ion interactions of S but 37% weaker ion−water interaction potential, resulting in a 40% reduction in the hydration enthalpy.71 The use of coarse-grained models based on short ranged interactions to study the thermodynamics of ion association in water is justified by the effective short-ranged attraction of ions in water. Due to the high dielectric constant of water, effective ion−ion interactions typically decay over 0.75 nm70,72 (the Bjerrum length of monovalent ions in water), irrespective of the sign of the charges in the ions.73 Glyoxal is represented by a united atom model (i.e., with all atoms except hydrogens), based on short-ranged anisotropic interactions and compatible with mW water. The parametrization of the glyoxal model is presented in section A of the Supporting Information. We do not add oxygen or nitrogen molecules to model the air interface, because 10096


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Journal of the American Chemical Society Figure 1. continued

Nliquid increasing linearly with ionic density, with a slope equal to 8.7, the average number of water molecules in the hydration shell of the ions.

ice surface contributes 2π nm2 of ice−liquid and liquid−vapor surfaces. As the ions approach, their solvation shells overlap, decreasing the area of the ice−liquid and liquid−vapor interfaces by ΔA. We compute the areas of the overlapped disks as a function of the distances between the ions (their centers) analytically, to determine ΔA(r) = sin θ − θ, where the angle in radians is θ = 2 cos−1(r/2). The change in length of the ice−liquid−vapor three phase interface, Δl(r), is estimated geometrically as the difference in perimeter between the overlapped and two separate circles of radius 1 nm, 4π(1 − θ/2π). We compute the diffusion coefficients of single S ion or glyoxal at 260 K in bulk liquid water and at the surface of ice. The self-diffusion coefficients D are determined from the Einstein relation,77 2dDt = limt→∞⟨|ri(t) − ri(0)|2⟩, where ⟨|ri(t) − ri(0)|2⟩ is the mean squared displacement and d is the dimensionality of the space in which the particles move: d = 2 at the ice surface and 3 in bulk. In the twodimensional case, we consider only contributions to the displacement in the two directions of the surface plane. The diffusion coefficient is calculated for a single ion and glyoxal, at the ice surface or in bulk, averaging over 100 ns simulations, with configurations saved every 20 ps. For the glyoxal, we considered the mean squared displacement of the center of the carbon−carbon bond.

3. RESULTS AND DISCUSSION 3.1. Ice−Air Interface Slows Down Ions and Speeds Up Organics with Respect to Liquid Water. We start by investigating the dynamics of a single ion at the ice surface. A previous simulation reported an 8-fold reduction in diffusivity of ions at the ice surface relative to bulk at temperatures 29 K below the melting point of the six-site water model.78,79 We compute the diffusion coefficient of a single S ion at the ice surface at 260 K, 13 K below the melting point of the mW water model.53 The disordered layer on ice is sparse at this temperature (Figure 1), exhibiting incomplete surface coverage (38 ± 1% of the surface is covered by liquid-like water). The disordered water is dynamic, grazing the entire ice surface by continuous melting and recrystallization of the ice interface. The simulations indicate that at 260 K the S ion diffuses continuously on the ice surface, with a diffusion coefficient that is 3 times slower than in bulk liquid water at the same temperature: Dion = 1.1 × 10−5 cm2 s−1 at the surface vs 3.3 × 10−5 cm2 s−1 in bulk water. The slowdown in diffusivity at the ice surface is due to the continuous melting and recrystallization of the hydration shell surrounding the ion as it moves through the ice surface. Our results indicate that, despite the sparse coverage of the surface by liquid-like water, ions at the ice surface at 260 K retain significant mobility. Different from the ions, we find that the mobility of glyoxal at the ice surface is about 25 times larger than in liquid water: at 260 K, Dglyoxal = 33 × 10−5 cm2 s−1 at the surface vs 1.3 × 10−5 cm2 s−1 in bulk water. While the diffusivities of glyoxal and the ion in liquid water are comparable, the diffusivities at the ice surface are very distinct. Glyoxal glides over the crystal surface, without experiencing the friction that arises from the continuous melting and reconstruction of the ice in its path. Sections 3.2 and 3.4 below show that the difference in mobility of ions and hydrophilic organics at the ice surface is due to their distinct abilities to increase the liquidity of the ice surface.

Figure 1. Ions increase the liquidity of the ice−air interface. Upper panel: Snapshots of the ice surface with increasing ion density. Ice is represented with silver sticks, liquid-like water molecules at the interfacial layer with blue balls, and ions with red balls. The snapshots A, B, C, and D correspond to surface ion densities 0, 0.05, 0.25, and 0.35 nm−2, respectively, and reveal the propensity of the ions to aggregate as solvent-separated clusters with increasing density of ions at the ice surface. The width of the disordered layer can reach up to ∼1 nm in the regions rich in ions. Lower panel: Number density of liquidlike water molecules at the ice surface (Nliquid) as a function of ionic surface density at 260 K. The clustering of ions at the ice surface is driven by the need to share the hydration shell water that reduces the amount of liquid water required at the ice surface to hydrate ions. This is reflected in the sublinear evolution (solid line) in the number of liquid molecules (Nliquid) at the ice surface with ion density. The dotted line represents the hypothetical scenario in which the ions at the ice surface do not share their hydration shells. This would result in 10097


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Journal of the American Chemical Society 3.2. Ions Experience Water-Driven Attraction at the Ice−Vapor Interface. Our simulations indicate that ions increase the amount of liquid-like water at the ice surface (lower panel of Figure 1), in agreement with the experiments.20,24 The simulations allow for a quantification of this effect. A single ion at the ice surface has 7.6 water molecules in the first hydration shell, about 10% less than the 8.6 found for its hydration in bulk liquid water. We find that, on average, a single S ion increases the liquidity of the ice surface by 8.7 water molecules (computed from the slope of Nliquid vs ion density at low surface concentrations, lower panel of Figure 1). The initial linear increase in liquidity with surface density of ions indicates that the interfacial ions build their solvation shell by locally melting ice. The amount of liquid-like water at the surface increases sublinearly for surface densities of ions larger than 0.05 nm−2 (lower panel of Figure 1). This is due to aggregation of solventseparated ions. The clustering behavior of ions at the ice−vapor interface points to a stronger effective attraction between ions at the ice surface, as there is no clustering of these ions in bulk liquid water (Supporting Information, Figure S3). The same, albeit less pronounced, effect is observed for ions that interact 40% less strongly with water (see Supporting Information C). Based on the hydration free energies of LiCl and NaCl,80 the latter should have properties close to those of S than to the weaker ions and produce a significant increase in the liquidity of the ice surface. In what follows we quantify the free energy of attraction between S ions at the ice−vapor interface and elucidate the origin of the underlying thermodynamic forces that drive their aggregation. To quantify the effect of the ice surface on the interaction between ions, we compute the free energy of attraction between a pair of ions at the ice surface, ΔGsurface(r), and compare it to the free energy of attraction of the same solutes in bulk liquid water (Figure 2). In both cases the ΔGsurface(r) shows a minimum at the solvent-separated distance for the ions. The attraction between ions at the ice surface at 260 K is 3.5 times stronger than that in bulk water, which explains the clustering of ions seen in Figure 1. The ion−ion attraction at the ice surface is not only stronger than that in bulk liquid water but also longer ranged, persisting up to ∼1.9 nm. This lengthscale is over three times larger than the effective range of attraction for the ions in bulk liquid water (Figure 2). This implies that the long-range attraction between ions at the ice surface is water-driven. To understand the factors that contribute to the attraction between ions at the ice−vapor interface we decompose the free energy of attraction between ions at the surface, ΔGsurface(r), into three contributions: (i) the direct free energy of association of ions embedded in liquid water, ΔGdirect(r), (ii) the cost of melting ice to create liquid water for the hydration of the ions, ΔGmelt(r), and (iii) the difference in surface free energy between the “liquid pool” of water created by the ions and the ice and vapor, ΔGint(r): ΔGsurface(r) = ΔGdirect(r) + ΔGmelt(r) + ΔGint(r). We approximate ΔGdirect(r) by the free energy of attraction in the ions in bulk water, because the number of water molecules in the first hydration shell of the ion is almost the same in bulk water and at the ice surface. ΔGdirect(r) is short-ranged and contributes ∼30% of the attraction between the ions at the surface. This implies that the surface specific water-driven contributions, ΔGmelt(r) and ΔGint(r), account for most of the driving force for ion−ion attraction at the ice surface.

Figure 2. Free energy of attraction between a pair of S ions at the ice surface and in bulk liquid water. ΔGsurface(r) is shown with full lines and the bulk ΔGbulk(r) with dashed lines. Blue indicates that the temperature is 260 K and red that is 270 K. The free energy of attraction between the S ions in bulk water displays a solvent separated pair (SSP) at 4.7 Å, same as predicted for LiCl in atomistic simulations with long-range electrostatic interactions.70 The attraction at the SSP is 1.2 kJmol−1 for the coarse-grained model and 3 kJ mol−1 for the atomistic model.70 The contact pair (CP) was not sampled in the simulations used to build this free energy profile and represents only 3% of the population of the atomistic model (separated from the SSP by a 22 kJ mol−1 barrier).70 The ion−ion interaction potential is zero at r > 4.3 Å. The free energy of ion−ion attraction at the ice surface highlights the increase in effective attraction compared to that in bulk. Additionally, the free energy of ion−ion attraction displays a characteristic long-range attraction different from that in bulk. The long-range attraction of ions at the ice surface is driven by the need to share the hydration shell water. This is reflected in the decrease in effective ion−ion attraction at 270 K as the enhanced liquidity at the ice surface reduces the necessity to share the hydration shell water.

The free energy contribution due to ice melting can be computed as ΔGmelt(r) = ΔμLSΔN(r), where ΔμLS is the difference in free energy between bulk ice and liquid at the corresponding temperature (which for mW water51,58 is in quantitative agreement with the experiment81) and ΔN(r) is the difference in the number of liquid-like water molecules when the ion pair is brought from “infinity” to a distance r. Figure 3a shows the spatial distribution of liquid-like water molecules around an ion pair for distances ranging from 0.5 to 4 nm. Each ion perturbs the liquidity on the ice surface at distances up to 1 nm. We note that the 1 nm shell around the ion consists of sparsely populated liquid-like molecules, not a thin film of liquid. As the ions come closer than 2 nm, they share liquid-like water molecules in their hydration shells (Figure 3b). The sharing of the molecules results in an attractive contribution to ΔGmelt(r) (Figure 3c) that accounts for ∼20% of the free energy of attraction of the solventseparated pair. The overlap in the hydration shells also decreases the area of ice−liquid and liquid−vapor interfaces and the ice−liquid−vapor contact line, which results in an interfacial contribution to the free energy of attraction, ΔGint(r) (see Supporting Information D), that decays over the same length scales than ΔGmelt(r) and should contribute the remaining 50% of the free energy of formation of the solvent-separated ion pair at the ice surface. It is important to note that the 2 nm length-scale of the free energy of ion−ion attraction at the ice surface (Figure 2) is the same as the one for ΔGmelt(r) (Figure 3c) and ΔGint(r) (Supporting Information, 10098


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indicate that a competition between an unfavorable entropic term, −TΔS = 26 kJ mol−1, and a favorable enthalpic contribution, ΔH = −30.5 kJ mol−1, is responsible for the favorable free energy of attraction, ΔGsurface = −4.5 kJ mol−1, between ions at the surface. 3.3. Attraction between Ions at the Ice Surface Is Anticooperative. The preceding analysis focused on the free energy of interaction for a pair of ions at the ice surface. To understand the behavior of a collection of ions at the ice surface and how they nucleate a solution phase, we must examine the free energy of attraction between multiple ions. Although it is practically impossible to compute the multidimensional free energy of attraction between a large collection of ions, we can determine whether the water-drive attraction between ions is cooperative, additive, or anticooperative. If the attraction between pairs of ions were additive, the multidimensional surface would be a linear combination of pair contributions, n−1 n ΔGsurface(r12, r13,... r(n‑1)n) = ∑i ∑i > j ΔGsurface(rij). If, instead, the attraction for the multiple ions were weaker (stronger) than the one predicted by this assumption, the attraction would be anticooperative (cooperative). To determine whether the clustering of ions shown in Figure 1b−d arises from cooperative, additive, or anticooperative interactions, we parametrize an effective ion−ion potential at the ice surface that exactly reproduces ΔGsurface(r) of a pair of ions at the surface at 260 K (see Supporting Information E). Using this potential, we evolve solvent-free two-dimensional Langevin simulations at 260 K with the same area and ion density of Figure 1d. We find that the ions cluster more aggressively in the implicit solvent than in explicit water (Supporting Information, Figure S3). The average largest cluster contains 22 ± 6 ions in the simulations with explicit water of Figure 1d and all 35 ions at the surface in the implicit solvent simulations. This implies that the free energy of attraction of multiple ions is less than the sum of the free energy of attraction of all of the pairs, i.e., the water-driven attraction of ions at the ice surface is anticooperative. We attribute this to a high degree of sharing of hydration shell waters per ion in a pair compared to that in a cluster. To illustrate this, consider an idealized case in which three ions, each with Nw number of water molecules in its solvation shell, are brought together from infinite distances to complete overlap of their solvation shells. In that case, the ΔN = 2Nw, but the pair additive model would predict 3Nw. We conclude that the smaller gain due to overlap between hydration shells of multiple ions is responsible for the anticooperativity for attraction between ions at the ice surface. 3.4. Organic Solutes do not Experience Water-Driven Attraction at the Ice Surface. Glyoxal (H2C2O2), the smallest dicarbonyl compound, is one of the most abundant oxygenated organics in the atmosphere.46−49 An estimated 45 teragrams per year of glyoxal are released to or produced in the atmosphere, where glyoxal participates in several chemical reactions and the nucleation of secondary organic aerosols.49 We investigate the effect of glyoxal on the ice surface at 260 K. Figure 4 shows that the free energy of attraction between a pair of glyoxal molecules on pristine ice is almost the same as for the two molecules in vacuum (4.8 vs 5.2 kJ mol−1) and much more pronounced than in liquid water (1.5 kJ mol−1). These results are consistent with the little effect that glyoxal has on the liquidity of the ice surface (Figure 5). At infinite surface dilution, each glyoxal molecule increases the liquidity of the ice

Figure 3. Effect of a pair of ions on the density of liquid-like water at the ice surface. (A) Degree of overlap between the hydration shells of an ion pair at the ice surface. The snapshots display the overlap of the time-averaged shell of hydration of an ion at the ice surface as the ion− ion distance is increased from 5 to 40 Å. (B) The number of water molecules shared per ion of the hydration shell water, ΔN(r), decreases as the two ions are pulled further apart. (C) The contribution due to the free energy of ice melting at 260 K, ΔGmelt(r), concurrently decreases with ΔN(r).

Figure S6). This indicates that the long-ranged length scale of the attraction between ions at the ice surface originates in the decrease in liquidity that results when ions share disordered water molecules in their hydration shells. The decrease in liquidity upon association is the driving force for ion−ion attraction at the ice surface. The thermodynamic driving force that brings the ions together decreases on heating from 260 to 270 K (Figure 2), as liquid-like water becomes less scarce5 at the ice surface. The direct attraction between ions, ΔGdirect(r) = ΔGbulk(r), does not change from 260 to 270 K (Figure 2). This indicates that the temperature dependence originates from the water driven components of the ion−ion interaction. We estimate the entropy of formation of the solvent separated pair (ssp) of ions from the change in free energy with temperature between 260 and 270 K, ΔS = −Δ(ΔGsurface(rSSP))/ΔT = −100 J mol−1 K−1. The magnitude of the entropy loss on association is intermediate between what would be predicted at 265 K from the crystallization of three water molecules, 3 × 19 J K−1 mol−1,51 and the entropy associated with the return of three molecules from the surface to the bulk of the liquid phase, 3 × 68 J K−1 mol−1.82 The negative value is consistent with the decrease in disordered water upon ion association. Our results 10099


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Figure 4. Free energy of attraction between a pair of glyoxal molecules at 260 K in bulk liquid water (solid cyan), the gas phase (dashed blue), at the ice surface (solid red), and at the ice surface with 0.38 nm−2 of S ions (solid green). The free energy of attraction between two glyoxal molecules at the ice surface indicates a preference for forming a direct contact pair, without any liquid water mediated interaction, same as in gas phase. Typically one of the glyoxal molecules in the contact pair configuration at the ice surface points one oxygen to the ice, same as reported for acetone and acetaldehyde in atomistic simulations.37,38 Increase in liquidity at the ice surface due to presence of ions decreases the free energy of attraction of the contact pair.

surface by three water molecules. However, Figure 5 shows that, at densities larger than 0.16 nm−2, the glyoxal molecules cluster into dry aggregates and do not increase the liquidity further. The clustering of glyoxal is driven by direct interactions between the organic molecules (Figure 4). We conclude that, unlike ions, hydrophilic organics do not experience waterdriven attraction at the ice surface. The increase in surface liquidity due to ions results in an enhanced adsorption of organics.24 To understand whether the synergism extends to the aggregation of these species, we compare the average size of the largest aggregate of glyoxal at the surface of pristine ice and ice with 0.35 nm−2 density of ions (Supporting Information, Figure S8). We find that dry agglomerates of glyoxal have some degree of sharing of liquid-like water molecules with the hydrated ion cluster that couples their diffusion at the ice surface. Less than 0.3 water molecules per glyoxal are shared with the ion cluster. The presence of ions does not impact significantly the average size of the largest cluster of glyoxal at the ice surface: the clustering decreasing slightly, consistent with the effect of surface liquidity on the glyoxal−glyoxal association free energy (Figure 4). We conclude that while ions may have a strong effect on the adsorption of organics to the ice surface, they do not have a significant impact on their aggregation. However, the coupling of the ion and glyoxal clusters could lead to an increase in the rate of reactions that necessitate contact of organics and ions or organics and liquid water.

Figure 5. Effect of glyoxal and/or ions on the liquidity of the ice surface. Snapshots of the ice surface with increasing glyoxal density (upper panel). Ice is represented with silver sticks, liquid-like water molecules at the ice surface with blue balls, the carbon and oxygen sites of glyoxal (Gl) are represented with silver and green balls, respectively, and the S ions with red balls. Snapshots A and B correspond to surface glyoxal density of 0.16 and 0.36 nm−2, respectively. Snapshot C corresponds to a surface glyoxal density 0.36 nm−2 in and ionic density 0.35 nm−2. Snapshot D corresponds to the ice surface with ionic density 0.35 nm−2. Lower panel: Number of liquid-like molecules at the ice surface at 260 K, Nliquid, as a function of surface density of glyoxal. The green solid line corresponds to the liquidity induced by glyoxal in the absence of ions, and the blue line in the presence of 0.35 nm−2 surface density of ions. The letters in the lower panel refer to concentrations illustrated by snapshots in the upper panel.

4. CONCLUSIONS The surface of atmospheric ice crystals is a sink for volatile atmospheric trace gases8,21 and also accumulates salts and other solutes excluded from the crystallization of water.18 Ions and organics are extremely insoluble in ice crystals. This makes the surface of ice the only possible locus for solvation of these species. In this study we use large-scale molecular dynamics simulations to characterize the effect of ions and oxygenated

organics on the liquidity of the ice surface−vapor interface at temperatures relevant to the troposphere and the role surface liquidity plays on the interaction between solutes and their aggregation at the surface. In agreement with experiments,20,24 the simulations indicate that ions strongly increase the amount of liquid-like water at the ice surface. We find that the growth in liquidity is sublinear with increasing ion density, due to the clustering of solvent-separated ions. The underlying thermody10100


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aggregation into dry agglomerates at the ice surface, which can enhance the propagation of photocatalyzed radical reactions involving oxygenated organics.29,94−96 We conclude that the ice surface imparts distinct characteristics to the solvation, aggregation, and mobility of ions and organics. The magnitude of the attraction of solutes at the ice surface is controlled by the scarcity of liquid-like water in the premelted layer. For both ions and organics, solute−solute attraction decreases with increase liquidity, although the origin of this decrease is different. For the ions, it is due to diminished water-driven attraction as liquid-like water becomes less scarce. For the hydrophilic organics, it is driven by a weakening of direct interactions as the molecules become solvated in liquidlike water. Our study supports the growing consensus that the premelted surface of ice has unique properties with respect to liquid water,1,9,18,21 and its specific characteristics should be considered in the modeling of the interactions and kinetics of reactive and nonreactive uptake of trace gases at environmental ice surfaces.

namic driving force behind this clustering is the existence of a water-mediated attraction between ions, mostly driven by the minimization of the free energy cost of interfacial contributions arising from the melting of water by the ions at the ice surface. We conclude that the decrease in liquidity upon association is the driving force for the water-driven attraction between ions at the ice surface. The driving force for ion association is enthalpy-driven: on approaching the melting temperature, the disordered water becomes less scarce at the ice surface and the drive for ion association wanes. Within 1 K of the melting temperature, the premelted layer would completely cover the ice surface and become on average thicker than 1 nm.5,16 A previous study has shown that such a thick premelted layer has orientational order and solvation properties toward methane identical to those of bulk liquid water at the same temperature;5 under these conditions we predict that there would be no water-driven attraction between ions at the ice surface. The length scale of ion−ion attraction at the surface of ice is controlled by the length scale of interference of the solvation shells of the ions. This results in an attraction at the surface of ice that is over three times longer-ranged than the effective length scale of attraction between a pair of ions in liquid water. The attraction between multiple ions is anticooperative, because their sharing of water molecules diminishes the strength of the attraction compared to the sum of attraction between individual pairs. Our results suggest atmospheric trace gases that dissociate at the ice surface (e.g., HNO3, HCl, etc.)83−88 will experience long-range water-mediated attraction leading to clustering of solvent separated ions at ice surfaces. This may lead to rapid nucleation of a concentrated solution phase at high enough partial pressures of the trace gases and could result in acceleration of the rates of chemical reactions27,89 and modifications of their mechanistic pathways.90−93 We find that the driving force for solute attraction at the ice surface is determined by the strength of water−solute attraction: ions interact very strongly with water and experience a water-mediated attraction at the ice surface, while hydrophilic organics do not. Hydrophilic volatile oxygenated organic compounds (e.g., acetone, acetaldehyde, and glyoxal) are too weak to produce a significant increase in the liquidity of the ice surface.19,22 The simulations indicate that glyoxal molecules perturb the surface of ice weakly and do not experience waterdriven attraction at the ice surface. Glyoxal nevertheless clusters at the ice surface, in the form of dry agglomerates, driven by the direct interactions between the organic molecules. The same was found for the interaction of purely hydrophobic polyaromatic molecules at the ice surface.44,45 In contrast with the synergism previously reported for the adsorption of organics at the ice surface in the presence of HCl,24 we do not find a significant effect of ions in the aggregation of organics at the ice surface. The strength of water−solute attraction not only determines the driving force for the interaction between solutes but also controls the mobility of solutes at the ice−air interface. Ions move slower at the ice surface than in bulk liquid water, because at the surface they experience a friction that arises from the continuous melting and reconstruction of ice in their path. Glyoxal glides over the ice surface without reconstructing it. As a result, glyoxal experiences a 25-fold speed up in diffusivity at the ice surface compared to that in liquid water. The enhanced mobility of organics adsorbed on ice favors their rapid

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/jacs.7b05233. (A) Glyoxal force field, (B) distribution of the largest cluster size of ions, (C) effect of the strength of ion− water attraction on the free energy of ion−ion attraction, and (D) interfacial contributions to the free energy of ion−ion attraction at the ice surface. (PDF)

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AUTHOR INFORMATION

Corresponding Author

*[email protected] ORCID

Valeria Molinero: 0000-0002-8577-4675 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Science Foundation through Award CHE-1305427 “Center for Aerosol Impacts on Climate and the Environment” and The University of Utah through an Undergraduate Research Opportunity Fellowship. We thank the Center for High Performance Computing at The University of Utah for technical support and a grant of computer time.

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REFERENCES

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