Tarso B. Ledur Kist - Open and Toroidal Electrophoresis

The relative electric (static) permittivity of water is represented by the constant and has an important role in the Coulomb equation , equation 1.1. This constant is high for water and Table 1.1shows a list for comparison. Moreover, the average number of hydrogen bonds between water molecules in pure water is much larger than the average number of hydrogen bonds within hydrofluoridic acid or ammonia, due to the symmetry of the two sp 3bonding orbitals and the two sp 3non-bonding orbitals of water molecules. This maximizes the average number of intermolecular hydrogen bonds – up to two as negative charge donor and two as negative charge acceptor – totaling up to four hydrogen bonds per molecule. Therefore, pure water is liquid from 273.15 K up to 373.15 K at 1 atm. The electric force ( ) among ions within a liquid is given by:

(1.1)

where represents the electric permittivity of the vacuum, represents the permittivity of the liquid (solvent), is defined as the relative permittivity, and are the charges, and the distance from each other. Table 1.1shows the relative (static) permittivity of some solvents; note that water is among the highest.

Figure 1.1 Some neighbors of oxygen in the periodic table. The atomic number is at the upper left, the electronic configuration at the upper right, and the electronegativity [2] at the bottom left. Oxygen has the highest electronegativity of all, except for fluorine. The hydrides shown are: methane, ammonia, water, hydrogen fluoride, silicon hydride, phosphine, hydrogen sulfide, and hydrogen chloride. All are gases in the 0 to 100 C range and one atmosphere of pressure, except for water and hydrogen fluoride, which has a boiling point of 19.5 C.

Neutral and non-ionizable solutes that are able to form hydrogen bonds tend to be soluble in water. Of course, solubility depends on the size of the molecule, the number of hydrogen bonds each molecule forms with water, the interaction energies of the solute molecules with themselves, and the entropy changes involved. If the number of hydrogen bonds is small and the solute molecules are large compared to the water molecules, then the solute tends to be expelled out from the bulk (interior) of the water by the water molecules. An exact prediction of solubility is given by the Gibbs free energy change (see Section 1.1.8).

The introduction of ions into certain solvents makes some solvent molecules attach around these ions, producing one or more layers of solvent molecules. The number of layers mainly depends on the charge of the ions, while the number of water molecules per layer largely depends on the size of the ions. These layers are a result of the charge–permanent dipole interaction between the ions and the permanent dipoles of the solvent molecules. Since water is a small molecule with a large electric dipole moment, this interaction is very strong in water. Figure 1.2illustrates the solvation of an anion in water and Figure 1.3illustrates the solvation of a cation in water.

Table 1.1 Relative (static) permittivities of liquids at 20 C. Extracted from Wohlfart and Lechner (2008) [3].

Liquid
N-methylformamide	189.0
Formamide	111.0
N-ethylformamide	104.7 a)
Water	80.100
Propylene carbonate	66.14
Methanoic acid (Formic acid)	51.1
Dimethyl sulfoxide	47.24
Glycerol	46.53
Ethylene glycol	41.4
N,N-dimethylacetamide	39.0
N,N-dimethylformamide	38.25
Acetonitrile	36.64
1,3-propanediol	35.1
Methanol	33.0
Ethanolamine	31.94
N,N-diethylformamide	29.6
2-pyrrolidone	28.18
Diethanolamine	25.75
Ethanol	25.3
Propan-2-one (Acetone)	21.01
1-propanol	20.8
2-propanol	20.18
Allyl alcohol	19.7
1-butanol	17.84
2-butanol	17.26
Cyclohexanone	16.1
Cyclopentanone	13.58
Trifluoroacetic acid	8.42
Tetrahydrofuran	7.6
Ethanoic acid (Acetic acid)	6.20
Propanoic acid	3.44
Dimethyl carbonate	3.087
1,4-dioxane	2.2189
Cyclohexane	2.0243
Hexane	1.8865

a)Calculated considering the change of 0.45% per C observed with the formamide derivatives at 20 C.

Figure 1.2 Schematic representation of chloride anion solvation in water.

Figure 1.3 Solvation of the sodium cation in water.