Conductance weak electrolytes

Ionic conductors arise whenever there are mobile ions present. In electrolyte solutions, such ions are nonually fonued by the dissolution of an ionic solid. Provided the dissolution leads to the complete separation of the ionic components to fonu essentially independent anions and cations, the electrolyte is tenued strong. By contrast, weak electrolytes, such as organic carboxylic acids, are present mainly in the undissociated fonu in solution, with the total ionic concentration orders of magnitude lower than the fonual concentration of the solute. Ionic conductivity will be treated in some detail below, but we initially concentrate on the equilibrium stmcture of liquids and ionic solutions. 

The conductivity of a solution containing such molecular ions may be small compared with the value that would result from complete dissociation into atomic ions. In this way, in the absence of neutral molecules, we can have a weak electrolyte. The association constant for (29) has a value that is, of course, the reciprocal of the dissociation constant for the molecular ion (PbCl)+ the logarithms of the two equilibrium constants have the same numerical value, but opposite sign. 

Incomplete Dissociation into Free Ions. As is well known, there are many substances which behave as a strong electrolyte when dissolved in one solvent, but as a weak electrolyte when dissolved in another solvent. In any solvent the Debye-IIiickel-Onsager theory predicts how the ions of a solute should behave in an applied electric field, if the solute is completely dissociated into free ions. When we wish to survey the electrical conductivity of those solutes which (in certain solvents) behave as weak electrolytes, we have to ask, in each case, the question posed in Sec. 20 in this solution is it true that, at any moment, every ion responds to the applied electric field in the way predicted by the Debye-Hiickel theory, or does a certain fraction of the solute fail to respond to the field in this way In cases where it is true that, at any moment, a certain fraction of the solute fails to contribute to the conductivity, we have to ask the further question is this failure due to the presence of short-range forces of attraction, or can it be due merely to the presence of strong electrostatic forces ... 

The conductivity of a 0.1 M acetic acid solution is much lower, however, than that of a 0.1 M hydrogen chloride solution. This and other experiments show that only a small fraction of the dissolved acetic acid, CH3COOH, has formed ions. Such a substance that dissolves and dissociates to ions only to a limited extent is called a weak electrolyte. 

Fig. 11-1. A strong electrolyte solution conducts better than a weak electrolyte solution.

Now we can explain the low conductivity of pure water. Though water dissociates into ions, H+(aq) and OH (aq), it does so only to a very slight extent. At equilibrium, the ion concentia-tions are only IQ-7 M. Water is a weak electrolyte. 

For strong electrolytes the molar conductivity increases as the dilution is increased, but it appears to approach a limiting value known as the molar conductivity at infinite dilution. The quantity A00 can be determined by graphical extrapolation for dilute solutions of strong electrolytes. For weak electrolytes the extrapolation method cannot be used for the determination of Ax but it may be calculated from the molar conductivities at infinite dilution of the respective ions, use being made of the Law of Independent Migration of Ions . At infinite dilution the ions are independent of each other, and each contributes its part of the total conductivity, thus ... 

The measurements of a by means of the electrical conductivity show that the dilution law holds good for weak electrolytes (a small), but for strong electrolytes (a large) it fails utterly. This behaviour has given rise to a considerable amount of discussion, a critical review of which will be found in a paper by the author ( Ionic Equilibrium in Solutions of Electrolytes ) in the Trans. Chem. Soc., 97, 1158, 1910. It appears that in this... 

It behaves as a weak electrolyte in ethanolamine, but the conductivity is high (Ref 3)... 

The solute in an aqueous strong electrolyte solution is present as ions that can conduct electricity through the solvent. The solutes in nonelectrolyte solutions are present as molecules. Only a small fraction of the solute molecules in weak electrolyte solutions are present as ions. 

FIGURE 1.4 Pure water is a poor conductor of electricity, as shown by the very dim glow ot the bulb in the circuit on the left (a). However, when ions arc present, as in an electrolyte solution, the solution does conduct. The ability of the solution to conduct is low when the solute is a weak electrolyte (b) but significant when the solute is a strong electrolyte (c), even when the solute concentration is the same in each instance. 

In aqueous electrolyte solutions the molar conductivities of the electrolyte. A, and of individual ions, Xj, always increase with decreasing solute concentration [cf. Eq. (7.11) for solutions of weak electrolytes, and Eq. (7.14) for solutions of strong electrolytes]. In nonaqueous solutions even this rule fails, and in some cases maxima and minima appear in the plots of A vs. c (Eig. 8.1). This tendency becomes stronger in solvents with low permittivity. This anomalons behavior of the nonaqueous solutions can be explained in terms of the various equilibria for ionic association (ion pairs or triplets) and complex formation. It is for the same reason that concentration changes often cause a drastic change in transport numbers of individual ions, which in some cases even assume values less than zero or more than unity. 

The infrared spectra suggest bidentate dithiocarbamate coordination, the magnetic susceptibility is markedly field-dependent. Although the compound is monomeric in CHCI3, a molar conductance in nitromethane is found, which increases sharply on dilution, attributed to a weak electrolytic character. Accordingly, the authors formulate this compound as [Nb2(Et2r/tc)sBr2]Br. 

Some chlorides exist, however, whose conductances are intermediate between those of good conductors and of insulators. This implies that even molten electrolytes can be categorized as strong, medium, and weak electrolytes. 

It has been seen above that the value of A, extrapolated to zero concentration provides A0, the equivalent conductance at infinite dilution, for strong electrolytes, HC1 and KC1. A similar operation for the determination of A, for the weak electrolytes will just not hold simply because, as it has been seen, weak electrolytes feature the fact their Ac rise steeply at high dilutions. The experimental determinations become very uncertain in these situations. 

It may be added that Kohlrausch s law does not lead to any method of deducing the contributions of the individual ions. The immediate practical application of Kohlrausch s law of independent contributions of the ions at infinite dilution is a method for deducing the limiting equivalent conductance, A0, of weak electrolytes. This will be illustrated by taking a specific example of a weak electrolyte. 

In a weak electrolyte such as CH3COOH, the A values rise steeply with decreasing concentration because more of the electrolyte ionizes according to the principle of equilibrium, and ionization is complete at infinite dilution. The sharp rise in the A value at lower concentration occurs because of a sharp increase in the number of ions in solution. Kohlrausch s law may be used in the determination of A0 for acetic acid or any weak electrolyte. According to this law, A0 for acetic acid is the sum of the ionic conductivities of H+ and CHjCOCT at infinite dilution... 

A+ = N A0. Thus, the ionic conductance of an ion is obtained by multiplying the equivalent conductivity at infinite dilution of any strong electrolyte containing that ion by its transport number. In this manner the ionic mobilities of the two ions present in the weak electrolyte can be calculated, and finally its equivalent conductivity at infinite dilution can be calculated by summing these two values. 

Let the electrolysis of dilute sulfuric acid (so-called electrolysis of water) with a platinum cathode and a platinum anode be considered next. Pure water is a very weak electrolyte and consequently a very poor conductor of electricity. It dissociates very slightly into H+ ions (it may be recalled that in fact, H+ ions does not remain as such but forms hydronium in H30+ by combining with a molecule of water, H+ + H20 H30+) and OFT ions. In the presence of little sulfuric acid (or for that matter any other strong electrolyte) the conductivity, i.e., ionization is greatly increased. The acidified water now contains H+ ions, OFT and SC3 ions. During electrolysis with platinum electrodes, H+ ions are attracted to the cathode, where each ion gains an electron and becomes a hydrogen atom ... 

Arrhenius postulated in 1887 that an appreciable fraction of electrolyte in water dissociates to free ions, which are responsible for the electrical conductance of its aqueous solution. Later Kohlrausch plotted the equivalent conductivities of an electrolyte at a constant temperature against the square root of its concentration he found a slow linear increase of A with increasing dilution for so-called strong electrolytes (salts), but a tangential increase for weak electrolytes (weak acids and bases). Hence the equivalent conductivity of an electrolyte reaches a limiting value at infinite dilution, defined as... 

A study of the concentration dependence of the molar conductivity, carried out by a number of authors, primarily F. W. G. Kohlrausch and W. Ostwald, revealed that these dependences are of two types (see Fig. 2.5) and thus, apparently, there are two types of electrolytes. Those that are fully dissociated so that their molecules are not present in the solution are called strong electrolytes, while those that dissociate incompletely are weak electrolytes. Ions as well as molecules are present in solution of a weak electrolyte at finite dilution. However, this distinction is not very accurate as, at higher concentration, the strong electrolytes associate forming ion pairs (see Section 1.2.4). 

The first theory of solutions of weak electrolytes was formulated in 1887 by S. Arrhenius (see Section 1.1.4). If the molar conductivity is introduced into the equations following from Arrhenius concepts of weak electrolytes, Eq. (2.4.17) is obtained, known as the Ostwald dilution law this equation... 

This equation is valid for both strong and weak electrolytes, as a = 1 at the limiting dilution. The quantities A = zf- FU have the significance of ionic conductivities at infinite dilution. The Kohlrausch law of independent ionic conductivities holds for a solution containing an arbitrary number of ion species. At limiting dilution, all the ions conduct electric current independently the total conductivity of the solution is the sum of the contributions of the individual ions. 

Interionic forces are relatively less important for weak electrolytes because the concentrations of ions are relatively rather low as a result of incomplete dissociation. Thus, in agreement with the classical (Arrhenius) theory of weak electrolytes, the concentration dependence of the molar conductivity can be attributed approximately to the dependence of the degree of dissociation a on the concentration. If the degree of dissociation... 

Fig. 2.5 Dependence of the molar conductivity of the strong electrolyte (HC1) and of the weak electrolyte (CH3COOH) on the square root of concentration...

The conductivity also increases in solutions of weak electrolytes. This second Wien effect (or field dissociation effect) is a result of the effect of the electric field on the dissociation equilibria in weak electrolytes. For example, from a kinetic point of view, the equilibrium between a weak acid HA, its anion A" and the oxonium ion H30+ has a dynamic character ... 

Electrolytes are defined as substances whose aqueous solutions conduct electricity due to the presence of ions in solution. Acids, soluble bases and soluble salts are electrolytes. Measuring the extent to which a substance s aqueous solution conducts electricity is how chemists determine whether it is a strong or weak electrolyte. If the solution conducts electricity well, the solute is a strong electrolyte, like the strong acid, HC1 if it conducts electricity poorly, the solute is a weak electrolyte, like the weak acid, HF. 

To distinguish between strong electrolytes, weak electrolytes and nonelectrolytes, prepare equimolar aqueous solutions of the compounds and test their electrical conductivity. If a compound s solution conducts electricity well, it is a strong electrolyte if its solution conducts electricity poorly, it is a weak electrolyte. A solution of a nonelectrolyte does not conduct electricity at all. 

Dilute aqueous solutions of weak acids, such as HF and HN02, contain relatively few ions because they are weak electrolytes, ionize only slightly into their ions and are therefore, poor conductors of electricity. This can be demonstrated using a conductivity apparatus as shown in Figure 6-1. 

The electrical conductivity of BaS04 is closest to that of C6Hi206, an organic molecule, which does not dissociate this observation further supports the previous evidence of the weak-electrolyte properties of BaS04. 

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