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Figure 19 shows that small ions cause electrostatic ordering of solvating waters, while large ions cause hydrophobic ordering shown in the Mercedes-Benz-plus-dipole model , In contrast, the charge at the center of a large ion cannot come close to a water molecule, so the electrostatic interactions with water are weak.

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Around large ions, water molecules form hydrogen bonds with other water molecules, as they would do around nonpolar solutes. An anion does not have the same effects as a cation of the same size. This leads to a notable difference between anion and cation size required to achieve a given level of water ordering. At very low concentrations, nonpolar solutes will dissolve in water.

Increasing its concentration causes a solute to reach its solubility limit, and to then aggregate. Now, to these hydrophobe—water solutions, add salt. Some types of salts will increase the solubilities of nonpolar solutes called salting in , and other types of salts will decrease the solubilities called salting out. Small ions tend to cause salting out: adding salt reduces hydrophobic solubilities in water. Large complex ions tend to cause salting in: adding salt increases nonpolar solubilities in water. Molecular dynamics simulations of Smith and Kalra et al.

Figure 20 illustrates the structural basis for Hofmeister effects. This exclusion increases the hydrophobe concentration in the remaining volume of the solution, driving the hydrophobes to aggregate with increasing ion concentration. Large ions do not exclude hydrophobes because large ions do not bind waters so tightly. Therefore, adding large ions to solution does not drive increased hydrophobe concentration in the remaining solution volume. The black bars in Figure 20 show the hydrophobe exclusion volume of different ions, as well as the well-known observation that the smaller ions tend to have the larger exclusion volumes.

Figure 21 shows the PMF of two ions coming together in water. It shows how the solvation-shell waters are structured at different ion separations. When the two ions are separated by about one layer of water, the bridging waters between them will be structured by multiple interactions. Each bridging water interacts with other bridging waters through hydrogen bonding, and each bridging water interacts with each ion through its water dipole.

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When the two mobile ions come into contact, the ion—ion electrostatics can also contribute substantially to the free energy. Extensive computer simulations using different water models show that the shape of the PMF depends on all these factors. The resulting free energy from the sum of the factors can be quite different for ions of different sizes and shapes. There is an interesting puzzle of ion pairing. Some salts are more soluble in water than others.

When a salt is composed of a small anion and a small cation, say LiF, it is relatively insoluble in water. When a salt is composed of a big anion and big cation, say CsI, it is also relatively insoluble in water. But, when a salt is composed of a small ion and a large ion, say CsF, then it is relatively soluble.

Collins explained this through his law of matching water affinities. Two small ions stick together their contact state is most stable because their ion—ion charge attractions dominate the energetics. Two large ions stick together because they act like hydrophobes since their charge interactions are weak, because the two ions cannot come sterically close together. The reason this state is stable is because the small ion attracts a water cage around it for electrostatic reasons, and the large ion is compatible with a water cage for hydrophobic reasons.

Therefore, the properties of aqueous solutions of even the simplest ions result from a subtle balance of geometry, hydrogen bonding, and charge interactions.

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When additional forces are also involved, such as when ions are near curved surfaces or protein binding sites, it can further tip that balance. In these cases, charge and geometry are less complicating factors. Water is commonly in contact with surfaces. Examples include water permeating through granular or porous or supramolecular structures or gels, or inside crowded biological cells, or bound to proteins or DNA, or at interfaces with air or oils.

For example, since water cannot form hydrogen bonds to hydrophobic surfaces, water tends to move away from them toward locations most compatible with forming water—water hydrogen bonding instead. At surfaces with air, water can lose hydrogen bonds simply because of the geometric constraint imposed by the surface; , see Figure Hydrophobes in water tend to concentrate at air—water interfaces.

This is because hydrophobes tend to localize wherever they are able to break the fewest water—water hydrogen bonds. Surfaces have a lower density of water—water hydrogen bonds than bulk water has. For example, consider acetic acid in a solution with water at a bulk pH equal to the p K a i. In the bulk, there will be equal concentrations of the protonated and deprotonated forms. Water inside nanotubes can be sterically constrained; see Figure In narrow hydrophobic pores Figure 24 , left , waters can form hydrogen bonded chains, where each water hydrogen bonds to one water neighbor in front and one behind.

According to semiempirical simulations by Hummer et al. Its high mobility in nanotubes is because water does not break or make hydrogen bonds to the nanotube walls as it flows. However, other studies show the opposite; namely that water flow is retarded in certain kind of nanotubes. Depending on the diameter, as well as their interior surface characteristics, nanotubes may exhibit ion selectivity, similar to that of the ion channels, which is attributed mostly to the formation of ion hydration shells.

What happens if water is squeezed in only a single direction, such as between two smooth plates? Sometimes ice itself is the constraining vessel. Other molecules can be captured within ices, in the form of clathrate hydrates.

Clathrate hydrates are ice cages that encapsulate small molecules, typically hydrocarbons. There are several regular clathrate structures, and Figure 25 shows the two most common forms. The shape and size of the caged solutes, as well as the pressure and temperature, will direct the type of clathrate structure adopted by the surrounding water.

Among the most notable clathrate hydrates is methane clathrate. Clathrate hydrates can be problematic commercially because they limit gas and oil extraction by clogging up the transport of hydrocarbons through pipelines. Despite this troublesome aspect, their large natural abundance and high methane density mean methane clathrates represent an enormous potential energy source.

1 Introduction: Why Is Water So Important?

Under the proper conditions, clathrate hydrates can be grown, often by doping ice using a hydrogen-bonding solute that can catalyze hydrate formation through the uptake of surrounding gas solutes. Growing such clathrate hydrates also provides an avenue for finding new structures of pure-water ices. Solutes can be extracted from clathrate hydrates by vacuum evacution to empty the ice cages, leading to new potential phases of pure ice.

Not unexpectedly, it is more delicate than the unevacuated Ne clathrate, collapsing at temperatures above K. Cooling a liquid slowly causes it to freeze to a stable crystalline state. But cooling a liquid rapidly can lead to kinetically trapped states that are glassy , i.

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In water, the characteristic relaxation times become of the order of s and the rate of change of the volume and entropy decreases abruptly but continuously to that of a crystalline solid. But, such simulations have been challenging because the time scales are very slow—15 orders of magnitude slower than in normal liquid water.

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But, even as glassy materials go, water is unusual. He defines fragile glass formers as having curved lines, where the viscosity is insensitive to temperature in the hot liquid and strongly dependent on temperature in the cold liquid. Strong glass formers are regarded as having memory of their molecular arrangements above their glass transition. Fragile liquids forget their glass structure much faster upon crossing the glass transition into the liquid phase. Water is more complicated than either of these behaviors, because of the complexity of structure in the supercooled liquid state.

The high-density and low-density components of water each have different viscosity characteristics. In simple liquids and in hot water , increasing the pressure increases the viscosity and decreases the diffusivity because applying pressure leads to crowding of the molecules, making their motion more sluggish. However, cold water behaves differently.

At temperatures below K the viscosity of water decreases with increasing pressure. In the case of cold water, applying pressure shifts water structures, crunching cage-like waters into van der Waals clusters. This breakage of hydrogen bonding frees up waters to move faster. Small molecules and ions diffuse through liquids. Their diffusion rates typically depend on the radius of the diffusant and the viscosity of the solvent. However, when the liquid is water, there is a remarkable exception. Ionic diffusion in water plays crucial roles in processes in biology e.