Solvents The use of solvents for analytical work is determined by their properties, as shown in Table 1.
Solvents with high dielectric constants (er > 10), for example, water and ammonia, are referred to as polar and are ionizing solvents, promoting the formation and separation of ions in their solutions, whereas those where eris about 2, such as diethyl ether, tetrachloromethane and hexane are nonpolar and are nonionizing solvents. There are also many solvents whose behavior is inter-mediate between these extremes.
The solution process in a liquid may be represented by a general equation:
+ =
The action of solution changes the properties of both solute and solvent. The solute is made more mobile in solution, and its species may solvate by attraction
B(sol)
The major component of a solution is referred to as the solvent, and there is a wide range of inorganic and organic solvents used in analytical chemistry. Their properties determine their use.
When a substance called the solute is dissolved in a solvent to form a solution, its behavior is often altered. Reactions in solution are faster than in the solid state. The amount of substance that can dissolve in a given amount of solvent at a stated temperature and pressure is called the solubility and is determined by the nature of the materials and the laws governing the solubility equilibrium.
Some substances form ions, which are species possessing a charge. These behave in a distinct way in solution. They may attract molecules of solvent, may associate together, and may react with other species to form complexes or a precipitate.
Since concentrations vary over a very wide range, they are often
represented by the logarithmic pX notation where pX = - log(X), where X is the concentration or activity of an ion, or an equilibrium constant.
The laws of thermodynamics govern the behavior of all species in solution. Every reaction depends upon the thermodynamic properties of the species involved. Where those properties are changed by the solvent by association, by reaction or temperature, the behavior will alter.
Physical and chemical equilibria in solution are most important.
Related topics Other topics in Section C Separation techniques (D1-D9) (C2-C10)
to the solvent. The solvent structure is also disrupted by the presence of species different in size, shape and polarity from the solvent molecules.
Ideally, the behavior should depend on the concentration m (in molarity, mole fraction or other units), but often this must be modified and the activity, a, used:
a = m g = p/pn
where g is called the activity coefficient. The vapor pressure of the solution is p, and that in the standard state is pn. Activities are dimensionless.
Solvents, such as water, with high dielectric constants (or relative permittivi-ties) reduce the force F between ions of charges z1e and z2e a distance r apart:
F = z1z2e2/eoerr2
where eo is the permittivity of free space. Also, they will solvate ions more strongly and thus assist ion formation and separation.
Hexane, diethyl ether and tetrachloromethane (CCl4) all have low dielectric constants and are nonpolar. They are very poor at ionizing solutes. However, they are very good solvents for nonpolar substances.
Solubility The equilibrium amount of solute which will dissolve in a given amount of solvent at a given temperature and pressure is called the solubility. The solu-bility may be quoted in any units of concentration, for example, mol m-3, molarity, mole fraction, mass per unit volume or parts per million (ppm).
There is a general ‘rule of thumb’ that ‘like dissolves like’. For example, a nonpolar hydrocarbon solvent such as hexane would be a very good solvent for solid hydrocarbons such as dodecane or naphthalene. An ester would be a good solvent for esters, and water or other polar solvents are appropriate for polar and ionic compounds.
● Gases dissolve in solvents according to Henry’s Law, provided they do not react with the solvent:
pB= xBK
where xBis the mole fraction of solute gas B which dissolves at a partial pres-sure pBof B, and K is a constant at a given temperature. This is analytically important for several reasons. For example, nitrogen is bubbled through solutions to decrease the partial pressure of oxygen in electrochemical experi-ments. Similarly, air is removed from liquid chromatography solvents by Table 1. Properties of some solvents
Solvent Boiling point (∞C) Density, Dielectric
(g cm-3) constant, er
Water 100 1.00 78.6
Ammonia -34 0.68 22.0
Ethanol 78 0.79 24.3
n-hexane 69 0.66 1.88
Diethyl ether 34 0.71 4.33
Note: density at 25∞C or at BP; dielectric constant = relative permittivity
passing helium through them, or by boiling them, since gas solubility decreases as the temperature is increased.
● Liquids. When different liquids are mixed, many types of behavior may occur. If the molecules in the liquids are of similar size, shape, polarity and chemical nature they may mix in all proportions. For example, benzene and methylbenzene (toluene) mix completely. In such ideal solutions, obeying Raoult’s law, the activity coefficient is close to 1:
a = p/pn= x
If the component molecules differ greatly in polarity, size or chemical nature (e.g., water and tetrachloromethane) they may not mix at all. This is an important condition for solvent extraction (Topic D1). The distribution of a solute between a pair of immiscible liquids depends primarily on the solubility of the solute in each liquid.
● Solids generally follow the ‘like dissolves like’ rule. Nonpolar, covalent materials dissolve best in nonpolar solvents. Solid triglycerides such as tristearin are extracted by diethyl ether, but are nearly insoluble in water.
Salts, such as sodium chloride are highly soluble in water, but virtually insoluble in ether.
Ions in solution The behavior of ions in solution may be summarized as follows.
(i) Solids whose structure consists of ions held together by electrostatic forces (e.g. NaCl) must be separated into discrete ions when they dissolve. These ions often gain stability by solvation with molecules of the solvent. Such solutions are described as strong electrolytes.
(ii) Some covalent molecules, such as ethanoic acid, may form ions by the unequal breaking of a covalent bond, followed by stabilization by solvation.
This occurs only partially and these are called weak electrolytes.
H : OCOCH3[ H++-OCOCH3
(iii) In some cases ions do not separate completely. In concentrated solutions, oppositely charged ions may exist as ion-pairs, large ions of surfactants may aggregate into micelles, which are used in capillary electrophoresis (Topics D8 and D9), and the dissociation of covalent molecules may be only partial.
(iv) At ‘infinite dilution’, that is, as the concentration approaches zero, ions are truly separate. However, even at quite low concentrations, the attractions between ions of opposite charge will cause each ion to become surrounded by an irregular cloud, or ionic atmosphere. As the solution becomes even more concentrated, the ionic atmosphere becomes more compact around each ion and alters its behavior greatly.
In very dilute solutions, the effects of the ionic atmosphere may be approxi-mated by using the Debye-Hückel theory, which predicts that the mean ionic activity coefficient, g± is, for an electrolyte with positive ions with charge z+, negative ions with charge z-, is given by:
log(g±) = -A (z+. z-) √|(I)
where I is the ionic strength =1⁄2S (cizi2) for all the ions in the solution.
For more concentrated ionic solutions, above 0.1 M, no general theory exists,
C1 – Solution equilibria 57
but additional terms are often added to the Debye-Hückel equation to compen-sate for the change in the activity.
The pX notation The concentration of species in solution may range from very small to large. For example in a saturated aqueous solution of silver chloride, the concentration of silver ions is about 10-5 M, while for concentrated hydrochloric acid the concentration of hydrogen and chloride ions is about 10 M. For convenience, a logarithmic scale is often used:
pX = -log (X)
where X is the concentration of the species, or a related quantity. Thus, for the examples above, pAg = 5 in saturated aqueous silver chloride and pH = -1 in concentrated HCl.
Since equilibrium constants are derived from activities or concentrations as noted below, this notation is also used for them:
pK = -log (K)
Most reactions will eventually reach equilibrium. That is, the concentrations of reactants and products change no further, since the rates of the forward and reverse reactions are the same.
From the above arguments concerning solutions, and from the laws of thermodynamics, any equilibrium in solution involving species D, F, U and V:
D + F [ U + V
will have an equilibrium constant , KT, at a particular temperature T given by:
KT= (aU. aV)/(aD. aF)
where the activities are the values at equilibrium. It should be noted that KT
changes with temperature. The larger the equilibrium constant, the greater will be the ratio of products to reactants at equilibrium.
There are many types of equilibria that occur in solution, but for the impor-tant analytical conditions of ionic equilibria in aqueous solution, four examples are very important.
(i) Acid and base dissociation. In aqueous solution, strong electrolytes (e.g., NaCl, HNO3, NaOH) exist in their ionic forms all the time. However, weak electrolytes exhibit dissociation equilibria. For ethanoic acid, for example:
HOOCCH3+ H2O [ H3O++ CH3COO -Ka= (aH. aA)/(aHA. aW) = 1.75 ¥ 10-5
where HA, W, H and A represent each of the species in the above equilibrium. In dilute solutions the activity of the water aW is close to 1.
For ammonia:
NH3+ H2O [ NH4++ OH -Kb = (aNH4+. aOH-)/(aNH3. aW) = 1.76 ¥ 10-5 Waterbehaves in a similar way:
2 H2O [ H3O++ OH -KW= (aH3O+. aOH-) = 10-14 Equilibria in
solution
(ii) Complexation equilibria. The reaction between an acceptor metal ion M and a ligand L to form a complex ML is characterized by an equilibrium constant. This is discussed further in Topic C7, but a simple example will suffice here:
M(aq) + L(aq) [ ML(aq) Kf= (aML)/(aM. aL)
For example, for the copper-EDTA complex at 25oC: Kf= 6.3 ¥ 1018
(iii) Solubility equilibria. If a compound is practically insoluble in water, this is useful analytically because it provides a means of separating this compound from others that are soluble. The technique of gravimetric analysishas been developed to give very accurate analyses of materials by weighing pure precipitates of insoluble compounds to give quantitative measurements of their concentration. For the quantitative determination of sulfate ions, SO4
2-, the solution may be treated with a solution of a soluble barium salt such as barium chloride BaCl2, when the following reaction occurs:
Ba2++ SO42-[ BaSO4(s) Conversely, if solid barium sulfate is put into water:
BaSO4(s) = Ba2++ SO4
2-The solubility product, Ksp, is an equilibrium constant for this reaction Ksp = a(Ba2+) . a(SO42-) = 1.2 ¥ 10-10
bearing in mind that the pure, solid BaSO4 has a = 1. This means that a solution of barium sulfate in pure water has a concentration of sulfate ions of only 1.1 ¥ 10-5M. The concentration of the barium ions is the same.
(iv) Redox equilibria. When a species gains electrons during a reaction, it undergoes reduction and, conversely, when a species loses electrons it undergoes oxidation. In the total reaction, these processes occur simultane-ously, for example:
Ce4++ Fe2+= Ce3++ Fe3+
The cerium is reduced from oxidation state 4 to 3, while the iron is oxidized from 2 to 3. Any general ‘redox process’ may be written:
Ox1 + Red2 = Red1 + Ox2
The equilibrium constant of redox reactions is generally expressed in terms of the appropriate electrode potentials (Topics C5, C8), but for the above reaction:
K = (a(Ce3+) . a(Fe3+))/(a(Ce4+) . a(Fe2+)) = 2.2 ¥ 1012
Summary
For ionic equilibria in solution, which are widely used in analytical chemistry, a large equilibrium constant for the reaction indicates that it will proceed practically to completion. If the equilibrium constant is of the order of 1010, then the ratio of products to reactants will be much greater than 1000 to 1. For example:
C1 – Solution equilibria 59
H++ OH-= H2O K = 1014 C(aq) + A(aq) = CA(solid) K = 1010 M(aq) + L(aq) = ML(complex) K = 1010 Ox1 + Red2 = Red1 + Ox2 K = 1012
Therefore, these reactions may be used for quantitative measurements, for example by volumetric or gravimetric techniques (Topics C5, C7 and C8).
It should be noted that, in calculations involving solution equilibria, certain rules should always be considered.
● Electroneutrality. The concentrations of positive and negative charges must be equal. Sometimes, ions that do not react are omitted from the equations, although they must be present in the solution.
● Stoichiometry. The total amounts of all species containing an element must be constant, since no element can be created or destroyed.
● Equilibria. All possible equilibria, including those involving the solvent, must be taken into account.
Section C – Analytical reactions in solution