Electrolytes solutions, colligative properties

So far we have discussed the colligative properties of nonelectrolyte solutions. Because electrolytes undergo dissociation when dissolved in water [W Section 4.1], we must consider them separately. Recall, for example, that when NaCl dissolves in water, it dissociates into Na Co ) and C aq). For every mole of NaCl dissolved, we get two moles of ions in solution. Similarly, when a formula unit of CaCL dissolves, we get three ions one Ca ion and two Cl ions. Thus, for every mole of CaCl2 dissolved, we get three moles of ions in solution. Colligative properties depend only on the number of dissolved particle.s—not on the type of particles. This means that a 0.1 m solution of NaCl will exhibit a freezing point depression twice that of a 0.1 m solution of a nonelectrolyte, such as sucrose. Similarly, we expect a 0.1 m solution of CaCL to depress the freezing point of water three times as much as a 0.1 m sucrose solution. To account for this effect, we introduce and define a quantity called the van t Hoff factor (i), which is given by... 

Solutions of Electrolytes—Calculating colligative properties of electrolyte solutions is more difficult than for solutions of nonelectrolytes. The solute particles in electrolyte solutions are ions or ions and molecules. Calculations using equations (14.5) and (14.6) must be based on the total number of particles present, and the van t Hoff factor is introduced into these equations to reflect this number. In all but the most dilute solutions, composition must be in terms of activities— effective concentrations that take into accoimt interionic forces. 

As noted earlier, colligative properties of solutions are directly proportional to the concentration of solute particles. On this basis, it is reasonable to suppose that, at a given concentration, an electrolyte should have a greater effect on these properties than does a nonelectrolyte. When one mole of a nonelectrolyte such as glucose dissolves in water, one mole of solute molecules is obtained. On the other hand, one mole of the electrolyte NaCl yields two moles of ions (1 mol of Na+, 1 mol of Cl-). With CaCl three moles of ions are produced per mole of solute (1 mol of Ca2+, 2 mol of Cl-). 

In contrast to nonelectrolyte solutions, in the case of electrolyte solutions the col-ligative properties depart appreciably from the values following from the equations above, even in highly dilute electrolyte solutions that otherwise by all means can be regarded as ideal (anomalous colligative properties). 

A—Freezing-point depression is a colligative property, which depends on the number of particles present. The solution with the greatest concentration of particles will have the greatest depression. The concentration of particles in E (a non electrolyte) is 0.10 m. All other answers are strong electrolytes, and the concentration of particles in these may be calculated by multiplying the concentration by the van t Hoff factor. 

The second period, from 1890 to around 1920, was characterized by the idea of ionic dissociation and the equilibrium between neutral and ionic species. This model was used by Arrhenius to account for the concentration dependence of electrical conductivity and certain other properties of aqueous electrolytes. It was reinforced by the research of Van t Hoff on the colligative properties of solutions. However, the inability of ionic dissociation to explain quantitatively the properties of electrolyte solutions was soon recognized. 

According to modem theory, many strong electrolytes are completely dissociated in dilute solutions. The freezing-point lowering, however, does not indicate complete dissociation. For NaCl, the depression is not quite twice the amount calculated on the basis of the number of moles of NaCl added. In the solution, the ions attract one another to some extent therefore they do not behave as completely independent particles, as they would if they were nonelectrolytes. From the colligative properties, therefore, we can compute only the "apparent degree of dissociation" of a strong electrolyte in solution. 

C. Particles in solution determine colligative properties. The first three materials are strong electrolyte salts, and C2H6O2 is a non-electrolyte. 

There are many measurement techniques for activity coefficients. These include measuring the colligative property (osmotic coefficients) relationship, the junction potentials, the freezing point depression, or deviations from ideal solution theory of only one electrolyte. The osmotic coefficient method presented here can be used to determine activity coefficients of a 1 1 electrolyte in water. A vapor pressure osmometer (i.e., dew point osmometer) measures vapor pressure depression. 

These phenomena that were previously considered anomalies of the mentioned colligative properties of the solutions, have been dealt with by Arrhenius in his effort to explain such anomalies by his well known theory of electrolytic dissociation. According to this explanation the molecules of a dissolved electrolyte partly split to form smaller particles, i. e. ions, which from the thermodynamic point of view are as effective as the undissociated molecules themselves. As the number of particles of matter is thus greater, the manifestations of colligative properties are increased, compared to what they would be with an undissociated electrolyte. 

Unlike weak electrolytes, solutions of strong ones have a far higher specific conductance the rise of the latter with rising concentration is also much more rapid. There is another difference the anomalies ascertained in the colligative properties of strong electrolytes cannot be ascribed to partial dissociation of molecules to ions as in the case of weak electrolytes. 

In all other solutions the so called degree of dissociation, as determined from the measurement of some colligative property, merely indicates the magnitude of interionic forces, it cannot, however, be taken as a measure of the quantity of dissociated and undissociated molecules of the solute. A complete theory of strong electrolytes, at least of their diluted solutions, has been developed by Debye and Hiickel, this theory is the basis of modern electrochemistry. 

D) The van t Hoff factor is in the calculations for colligative properties of solutions. Because the number of solute particles in solution affects these factors, an adjustment must be made for electrolytic solutes. This is due to the fact that electrolytes, when dissolved, yield as many particles as the number of ions in the... 

Strong Electrolytes. Solutes of this type, such as HCl, are completely dissociated in ordinary dilute solutions. However, their colligative properties when interpreted in terms of ideal solutions appear to indicate that the dissociation is a little less than complete. This fact led Arrhenius to postulate that the dissociation of strong electrolytes is indeed incomplete. Subsequently this deviation in colligative behavior has been demonstrated to be an expected consequence of interionic attractions. 

The colligative properties of electrolyte solutions are described by including the van t Hoff factor in the appropriate equation. For example, for changes in freezing and boiling points the modified equation is... 

The theory of electrolytic dissociation, AVhereas the osmotic pressure and the other colligative properties of aqueous solutions of substances, such as cane sugar, obey van t Hoff s laws, marked deviations are met with in aqueous solutions of acids, bases, and salts, even at great dilutions. The osmotic pressure and lowering of the freezing point for these solutions are still found to be approximately proportional to the molecular concentration, but are considerably greater than the theoretical values. To allow for this van t Hoff introduced a new term into his osmotic pressure equation, writing for such solutions... 

COLLIGATIVE PROPERTIES OF ELECTROLYTES AS COMPARED WITH NON-ELECTROLYTE SOLUTES... 

On application of the van t Hoff equation to the dmg molecules in solution, consideration must be made of any ionisation of the molecules, since osmotic pressure, being a colligative property, will be dependent on the total number of particles in solution (including the free counterions). To allow for what was at the time considered to be anomalous behaviour of electrolyte solutions, van t Hoff introduced a correction factor, i. The value of this factor approaches a number equal to that of the number of ions, v, into which each molecule dissociates as the solution is progressively diluted. The ratio ijv is termed the practical osmotic coefficient, [Pg.69]

Nonelectrolytes in aqueous solution Many molecular compounds dissolve in solvents but do not ionize. Such solutions do not conduct an electric current, and the solutes are called nonelectrolytes. Sucrose is an example of a nonelectrolyte. A Im sucrose solution contains only one mole of sucrose particles. Figure 15-16 compares the conductivity of a solution containing an electrolyte solute with one containing a nonelectrolyte solute. Which compound would have the greater effect on colligative properties, sodium chloride or sucrose ... 

Acids are classified as strong or weak, depending on the extent to which they are ionized in solution. In a weak acid the transfer of hydrogen ions to water does not proceed to completion. A weak acid such as acetic acid is thus also a weak electrolyte its aqueous solutions do not conduct electricity as well as a strong acid of the same concentration because fewer ions are present. A weak acid shows smaller values for colligative properties than a strong acid (recall the effect of dissolved acetic acid on the freezing point of water in Fig. 11.13). 

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