Equipos de protección colectiva

Dipole-dipole interactions between polar solute particles in organic solvents

may be either repulsive (“head-head” interactions) or attractive (“head-tail” in- teractions). The former are rare and are due to steric hindrance around one of the poles of the dipole, which causes the ends of neighboring molecules that carry partial charges of the same sign to approach it (seeFig. 2.1b). The attrac- tive dipole interactions lead to solute aggregation, either to dimers, where the “head” of each of the two partners is near the “tail” of the other (see Fig. 2.1a), or to larger, chainlike or cyclic aggregates (oligomers), for which each member of the chain or cycle interacts with a neighbor at both ends (see Fig. 2.1c). The

Fig. 2.6 Ion pairs. A two-dimensional representation of (a) a solvent-separated pair of ions, each still retaining its intact shell of solvating solvent molecules; (b) a solvent- sharing ion pair, which has lost some of the solvent between the partners, so that one layer of solvent shared between them separates them; (c) a contact ion pair, the cation and anion being contiguous.

dipole moment, µ, is proportional to the spatial separation of the (partial) charges that cause the particle to be dipolar and is generally considered to “re- side” midway between these charges. The energy of dipole–dipole interactions depends on the product of the squares of the dipole moments, µ, of the partners (i.e., onµ2, if there is a single kind of interacting species) and inversely on the sixth power of the distance between the centers of the dipoles. Such an energy is not very large, ranging between 1 and 10 kJ mol−1 (i.e., 0.4 to 4 times the mean thermal energy, RT, that tends to disperse the aggregates). Tertiary amine salts in hydrocarbon solvents are typical examples of such aggregated solutes.

Hydrogen bonding occurring between solute particles also leads to aggre-

gation, when the molecule of the solute contains both a hydrogen atom bonded to an electronegative atom and another electronegative atom that can accept a hydrogen bond. Typical of such solutes are carboxylic acids and acidic phos-

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phate esters. Cyclic dimers are formed (seeFig. 2.1d), bonded by two hydrogen bridges, the energy of each being about 20 kJ mol−1. Such cyclic dimers are quite stable. If the numbers of ionizable hydrogen atoms and electronegative accepting sites do not match, then noncyclic aggregates, possibly larger than dimers, are formed by hydrogen bonding. One aspect of such an aggregation is a drastic increase in the viscosity of the solution with increasing concentra- tions.

The mean aggregation number of a solute at molality m, forming various oligomers (i-mers, i= 1, 2, 3, . . . ), is

%n m= /∑imi = ∑iimi/∑imi ( . )2 51

the molality of each oligomeric species being weighted by the number of mono- mers it contains. The equilibrium constants for the aggregation reactions can be obtained from the dependence of n˜ on m by the methods discussed elsewhere in this book (seeChapter 3).

The third kind of solute–solute interaction, donor–acceptor adduct forma-

tion, tends to form 1:1 adducts between the molecules of two different kinds of

solute, rather than to lead to self-aggregation. Adduct formation results when one partner has a donor atom (i.e., an atom with an exposed pair of unbonded electrons), such as the nitrogen atom in trioctylamine, and the other has an acceptor atom (i.e., one with an orbital that can take up such a pair of electrons), such as antimony in antimony pentachloride or even the hydroxyl hydrogen atom in octanol. Although adduct formation formally can be said to take place in dipole–dipole interactions between two different kinds of molecules, this is rare, since the strong dipoles required naturally participate also in self-aggrega- tion. But if one kind of molecule is a donor only and the other is an acceptor only, as occurs when octanol (the oxygen atom of which has two exposed non- bonded pairs of electrons) in the foregoing example is replaced by antimony pentachloride, then self-aggregation is precluded, and only mutual adduct for- mation can take place (see following).

Hydrogen bond formation between dissimilar molecules is an example of adduct formation, since the hydrogen atom that is bonded to an electronega- tive atom, such as oxygen or nitrogen, is a typical acceptor atom. The ability of molecules to donate a hydrogen bond is measured by their Taft–Kamlet

solvatochromic parameter,α, (or αm. for the monomer of self-associating sol- utes) (see Table 2.3). This is also a measure of their acidity (in the Lewis sense, see later, or the Brønsted sense, if protic). Acetic acid, for instance, hasα = 1.12, compared with 0.61 for phenol. However, this parameter is not necessarily correlated with the acid dissociation constant in aqueous solutions. The ability of molecules to accept a hydrogen bond is measured by the Taft– Kamlet solvatochromic parameter,β, (or βmfor the monomer of self-associat- ing solutes) (see Table 2.3). This, too, is a measure of their basicity (in the Lewis sense), also measured by the Gutmann donor number DN (discussed later). Thus, pyridine has β = 0.64, compared with 0.40 for acetonitrile, but

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again, this measure is not directly correlated to the base dissociation constant (protonation constant) in aqueous solutions. The tendency for a hydrogen bond to be formed between two dissimilar solute molecules A and B increases with the sum of the products of their donating and accepting parameters:αAβB+ αBβA.

Adduct formation does not require formation of a hydrogen bond, and other acceptor atoms (or molecules) are known (e.g., transition metal cations in general, SbCl5, I2, and so on). They are collectively called Lewis acids, and they

react with electron pair donors, that are collectively called Lewis bases.

The enthalpy change during donor–acceptor adduct formation has been re- lated by Drago to the sum of two terms: (1) the product of the electrostatic proper- ties of the acid and the base, EAand EB; and (2) the product of their tendency toward covalent bonding, CAand CB[10]. For the particular case, where the ac- ceptor is specified to be SbCl5(and the inert solvent is 1,2-dichloroethane), the negative of this enthalpy change (in kcal mol−1, 1 cal= 4.184 J) is the Gutmann donor number, DN [2,11]. These concepts are further discussed inChapter 3.

2.7

SOLUBILITIES IN BINARY SYSTEMS

Binary systems are systems that involve two components, that is, two substances

that can be added individually. (Ions are not components; they have to be added in combinations as electroneutral electrolytes; see section 2.5.) The phase rule (see section 2.2) states that at a given temperature and pressure, when only a single liquid phase is present, there will be one degree of freedom, and we can choose the composition of this phase, made up from the two components, at will. This does not preclude, however, the phenomenon of saturation, where beyond a certain amount of solute in the liquid mixture a new phase appears. When two phases are present at the given temperature and pressure in the binary system, there are no longer any degrees of freedom, and the compositions of the phases are fixed. The new phase may be a solid or a second liquid. Generally, the solid phase is the pure solid solute, but in rare circumstances, it is a solid solution of the two components in each other; occasionally, it is a pure solid solvate of the original solute by the solvent. When a second liquid phase separates out, it is generally a saturated solution itself, rather than a pure liquid. We then have a solvent-rich dilute solution of the solute and a solute-rich concentrated solution of it. As an example, consider a solution of phenol in water at 25°C. At this tempera- ture, pure phenol is a solid (its melting point is 40.9°C), but when equilibrated with water at 25°C, the saturated aqueous layer contains 8.66% by mass of phenol, and the phenol-rich layer contains 28.72% by mass of water. However, on a mole fraction basis, both layers appear to be water-rich phases, the one having xwater=

0.982 and the other (the “phenol” phase) having xwater= 0.678.

The solubility of a solute in a solvent is given by the composition of the

saturated solution at a given temperature and pressure. The solubility may be expressed on any of the concentration scales: the molar (mol L−1; i.e., per liter

of the saturated solution), the molal (mol kg−1; i.e., per kilogram of the pure solvent), the mole fraction, the mass fraction (wt%; seeTable 2.2), or the vol- ume fraction scales. Only if the solute is a gas is the pressure of any signifi- cance. For liquid and solid solutes, however, the temperature is the only variable that ordinarily needs to be specified.