Electrolyte, strong

Strong electrolytes are considered to be completely dissociated in their solutions and their deviations from ideal behavior are ascribed to the effect of the strong forces of attraction acting between the oppositely charged ions. As in such solutions there are no undissociated molecules apart from ions, no equilibrium is established to which the law of mass action could be applied. The potential of the electrolyte as a whole dissociated according to the formula 

The electrolyte in the standard state is also supposed to be fully dissociated, so that even in this case the potential p° of the electrolyte as a whole can be expressed as the sum of products of the number of gram-ions, formed on dissociation of one mole of the electrolyte and of their standard potentials  

On substituting into the equations (V-28) and (V-29) the relations between different potentials in accordance with (V-8) 

As the activity of individual ions cannot be determined separately, the con- 

If the concentration of an electrolyte in its standard state is expressed by its molality, the following equations can bo written for the activity coefficients in a concrete solution 

Arrhenius originally believed that his theory and the formulation of Ostwald would explain the conductivity behavior of both strong and weak electrolytes. However, it soon became apparent that strong electrolytes required another explanation. Several lines of evidence indicated this. One was that the plots of A against c for strong electrolytes, unlike those for weak ones, could not be fitted to Ostwald s equations. Another was that the heats of neutralization of solutions of strong acids and bases (e.g., the neutralization of HCl and NaOH) could be explained only if it was assumed that these electrolytes are completely dissociated over a considerable concentration range. 

The distribution of chloride ions round a sodium ion (a) in the crystal lattice, (b) in a solution of sodium chloride. In the solution the interionic distances are greater, and the distribution is not as regular, but near to the sodium ion there are more chloride ions than sodium ions, 

A second factor which retards the motion of an ion in solution is the tendency 

The second group of interactions affecting diffusion involves solute olvent interactions. In Section 6.3, we explore the extremely large solute-solvent interactions which occur near the spinodal limit, where phase separation is incipient. Diffusion in these regions leads to the phenomenon of spinodal decomposition, which is also discussed in Section 6.3. 

In the last section of this chapter, we summarize diffusion affected by solute-boundary interactions, which is the third important group of interactions. Solute-boundary interactions occur in porous solids with fluid-filled pores. They include such diverse phenomena as Knudsen diffusion, capillary condensation, and molecular sieving. Because these phenomena promise high selectivity for separations, they are an active area for research. They and the other interactions illustrate the chemical factors that can be hidden in the diffusion coefficients which are determined by experiment. 

Every high school chemistry student knows that when sodium chloride is dissolved in water, it is ionized. Sodium chloride in water does not diffuse as a single 

The diffusion of sodium chloride can be accurately described by a single diffusion coefficient. Somehow this does not seem surprising, because we always refer to sodium chloride as if it were a single solute and ignore the knowledge that it ionizes. We get away with this selective ignorance because the sodium and chloride ions diffuse at the same rate. If they did not do so, we could easily separate anions from cations. 

Another interesting result in Table 6.1-1 is that the sodium ion diffuses more slowly than the chloride ion. In other words, the sodium ion does not have the same diffusion 

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Kohlrausch equation This equation, which describes the behaviour of strong electrolytes on dilution, states that... 

Strong electrolytes are dissociated into ions that are also paired to some extent when tlie charges are high or the dielectric constant of the medium is low. We discuss their properties assuming that the ionized gas or solution is electrically neutral, i.e. 

The theory of strong electrolytes due to Debye and Htickel derives the exact limiting laws for low valence electrolytes and introduces the idea that the Coulomb interactions between ions are screened at finite ion concentrations. 

Salts of acids other than hydrochloric acid commonly show increased solubiUty in hydrochloric acid. This phenomenon has been explained by the Debye-Hbckel theory for strong electrolytes (17—19). 

Fig. 15. Ion movements in the electro dialysis process. Courtesy U.S. Agency for International Development, (a) Many of the substances which make up the total dissolved soHds in brackish water are strong electrolytes. When dissolved in water, they ionize ie, the compounds dissociate into ions which carry an electric charge. Typical of the ions in brackish water are Cl ,, HCO3, , and. These ions tend to attract the dipolar water molecules...

Battery electrolytes are concentrated solutions of strong electrolytes and the Debye-Huckel theory of dilute solutions is only an approximation. Typical values for the resistivity of battery electrolytes range from about 1 ohmcm for sulfuric acid [7664-93-9] H2SO4, in lead—acid batteries and for potassium hydroxide [1310-58-3] KOH, in alkaline cells to about 100 ohmcm for organic electrolytes in lithium [7439-93-2] Li, batteries. 

The apparent acid strength of boric acid is increased both by strong electrolytes that modify the stmcture and activity of the solvent water and by reagents that form complexes with B(OH) 4 and/or polyborate anions. More than one mechanism may be operative when salts of metal ions are involved. In the presence of excess calcium chloride the strength of boric acid becomes comparable to that of carboxyUc acids, and such solutions maybe titrated using strong base to a sharp phenolphthalein end point. Normally titrations of boric acid are carried out following addition of mannitol or sorbitol, which form stable chelate complexes with B(OH) 4 in a manner typical of polyhydroxy compounds. EquiUbria of the type ... 

Adsorption with strongly favorable isotherms and ion exchange between strong electrolytes can usually be carried out until most of the stoichiometric capacity of the sorbent has been utilized, corresponding to a thin MTZ. Consequently, the total capacity of the bed is... 

Viewing the dissociation of strong electrolytes another way, we see that the ions formed show little affinity for one another. For example, in HCl in water. Cl has very little affinity for H ... 

Sodium acetate, the sodium salt of acetic acid, is a strong electrolyte and dissociates completely in water to yield Na and Ac. )... 

Reliable pH data and activities of ions in strong electrolytes are not readily available. For this reason calculation of corrosion rate has been made using weight-loss data (of which a great deal is available in the literature) and concentration of the chemical in solution, expressed as a percentage on a weight of chemical/volume of solution basis. Because the concentration instead of the activity has been used, the equations are empirical nevertheless useful predictions of corrosion rate may be made using the equations. 

When an ionic solution contains neutral molecules, their presence may be inferred from the osmotic and thermodynamic properties of the solution. In addition there are two important effects that disclose the presence of neutral molecules (1) in many cases the absorption spectrum for visible or ultraviolet light is different for a neutral molecule in solution and for the ions into which it dissociates (2) historically, it has been mainly the electrical conductivity of solutions that has been studied to elucidate the relation between weak and strong electrolytes. For each ionic solution the conductivity problem may be stated as follows in this solution is it true that at any moment every ion responds to the applied field as a free ion, or must we say that a certain fraction of the solute fails to respond to the field as free ions, either because it consists of neutral undissociated molecules, or for some other reason ... 

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 ... 

A criterion for the presence of associated ion pairs was suggested by Bjerrum. This at first appeared to be somewhat arbitrary. An investigation by Fuoss,2 however, threw light on the details of the problem and set up a criterion that was the same as that suggested by Bjerrum. According to this criterion, atomic ions and small molecular ions will not behave as strong electrolytes in any solvent that has a dielectric constant less than about 40. Furthermore, di-divalent solutes will not behave as strong electrolytes even in aqueous solution.2 Both these predictions are borne out by the experimental data. 

When an ionic solid such as NaCl dissolves in water the solution formed contains Na+ and Cl- ions. Since ions are charged particles, the solution conducts an electric current (Figure 2.12) and we say that NaCl is a strong electrolyte. In contrast, a water solution of sugar, which is a molecular solid, does not conduct electricity. Sugar and other molecular solutes are nonelectrolytes. 

As we have pointed out, strong acids and bases are completely ionized in water. As a result, compounds such as HC1 and NaOH are strong electrolytes like NaCl. In contrast, molecular weak acids and weak bases are poor conductors because their water solutions contain relatively few ions. Hydrofluoric acid and ammonia are commonly described as weak electrolytes. 

The following figures represent species before and after they are dissolved in water. Classify each species as weak electrolyte, strong electrolyte, or nonelectrolyte. You may assume that species that dissociate during solution break up as ions. 

Strong electrolyte A compound that is completely ionized to ions in dilute water solution, 37 Strontium, 543 Strontium chromate, 434 Structural formula A formula showing the arrangement of bonded atoms in a molecule, 34,579-580,586,590,593, 597... 

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

In pure water, where the only source of ions is reaction (6), the concentrations of H+(aq) and OH (aq) must be equal. But what if we add some HC1 to the solution We have already noted that HQ is a strong electrolyte, dissolving to give the ions H+(aqJ and G (aq). Thus, hydrogen chloride adds H+(aq) but not OH (aq) to the solution. The concentrations [H+] and [OH-] are no longer equal. However, they are still found to be tied together by the equilibrium relationship... 

Suppose, on the other hand, that we add sodium hydroxide, NaOH, to pure water. Sodium hydroxide is also a strong electrolyte, adding... 

Sodium benzoate is a strong electrolyte its aqueous solutions contain sodium ions, Na+f ql, and benzoate ions, QHtCOO faqJ. Hence the equilibrium involved is the same as before ... 

It is important to realise that whilst complete dissociation occurs with strong electrolytes in aqueous solution, this does not mean that the effective concentrations of the ions are identical with their molar concentrations in any solution of the electrolyte if this were the case the variation of the osmotic properties of the solution with dilution could not be accounted for. The variation of colligative, e.g. osmotic, properties with dilution is ascribed to changes in the activity of the ions these are dependent upon the electrical forces between the ions. Expressions for the variations of the activity or of related quantities, applicable to dilute solutions, have also been deduced by the Debye-Hiickel theory. Further consideration of the concept of activity follows in Section 2.5. 

With an aqueous solution of a salt of class (1), neither do the anions have any tendency to combine with the hydrogen ions nor do the cations with the hydroxide ions of water, since the related acids and bases are strong electrolytes. The equilibrium between the hydrogen and hydroxide ions in water ... 

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