Electrolytes acids and bases

It should be noted that pKw is temperature-dependent and can affect the calculation. Table 2.7 presents the effect of pH on the ionization of weak electrolytes (acids and bases) and is taken from Doluisis and... 

Acids and Bases as Electrolytes Acids and bases are categorized in terms of their... 

Monovalent cations are good deflocculants for clay—water sHps and produce deflocculation by a cation exchange process, eg, Na" for Ca ". Low molecular weight polymer electrolytes and polyelectrolytes such as ammonium salts (see Ammonium compounds) are also good deflocculants for polar Hquids. Acids and bases can be used to control pH, surface charge, and the interparticle forces in most oxide ceramic—water suspensions. 

Hydrogen was recognized as the essential element in acids by H. Davy after his work on the hydrohalic acids, and theories of acids and bases have played an important role ever since. The electrolytic dissociation theory of S. A. Arrhenius and W. Ostwald in the 1880s, the introduction of the pH scale for hydrogen-ion concentrations by S. P. L. Sprensen in 1909, the theory of acid-base titrations and indicators, and J. N. Brdnsted s fruitful concept of acids and conjugate bases as proton donors and acceptors (1923) are other land marks (see p. 48). The di.scovery of ortho- and para-hydrogen in 1924, closely followed by the discovery of heavy hydrogen (deuterium) and... 

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. 

For very weak or slightly ionised electrolyes, the expression a2/( 1 — a) V = K reduces to a2 = KV or a = fKV, since a may be neglected in comparison with unity. Hence for any two weak acids or bases at a given dilution V (in L), we have a1 = y/K1 V and a2 = yjK2V, or ol1/ol2 = Jk1/ /K2. Expressed in words, for any two weak or slightly dissociated electrolytes at equal dilutions, the degrees of dissociation are proportional to the square roots of their ionisation constants. Some values for the dissociation constants at 25 °C for weak acids and bases are collected in Appendix 7. 

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 ... 

We saw in Section 1.1 that electrolytes are classified as strong or weak according to the extent to which they are present as ions in solution. We classify acids and bases in a similar fashion. In these definitions the term deprotonation means loss of a proton and protonation means the gain of a proton ... 

The first substantial constitutive concept of acid and bases came only in 1887 when Arrhenius applied the theory of electrolytic dissociation to acids and bases. An acid was defined as a substance that dissociated to hydrogen ions and anions in water (Day Selbin, 1969). For the first time, a base was defined in terms other than that of an antiacid and was regarded as a substance that dissociated in water into hydroxyl ions and cations. The reaction between an acid and a base was simply the combination of hydrogen and hydroxyl ions to form water. 

Walden, P. (1929). Salts, Acids and Bases Electrolytes, Stereochemistry. New York McGraw-Hill. 

It is a typical feature of aqueous electrolyte solutions that one can, within wide limits, change the solute concentrations and hence the conductivities themselves. Pure water has a very low value of o it is about 5 pS/m at room temperature after careful purification of the water. In the most highly conducting solutions (i.e., concentrated solutions of acids and bases), values of 80 S/m can be attained at the same temperature values seven orders of magnitude higher than those found for pure water. 

For symmetric electrolytes i=l for 1 2 electrolytes (e.g., Na2S04), 1 3 electrolytes (AICI3), and 2 3 electrolytes ([Al2(S04)3], the corresponding valnes of A, are 1.587, 2.280, and 2.551. Mean ionic activity coefficients of many salts, acids, and bases in binary aqneons solutions are reported for wide concentration ranges in special handbooks. 

The theory of electrolytic dissociation also provided the possibility for a transparent definition of the concept of acids and bases. According to the concepts of Arrhenius, an acid is a substance which upon dissociation forms hydrogen ions, and a base is a substance that forms hydroxyl ions. Later, these concepts were extended. 

When solutions of acids and bases are sufficiently dilute or when other electrolytes are present, the activity, rather the concentration of hydrogen ions, should be substituted in the pH equation. 

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... 

Acids and bases are almost always found as aqueous solutions— that is, dissolved in water. Solutions of both acids and bases are called electrolytes. Electrolytes conduct electricity, which is the movement of electrons or other charged particles. When an acid or a base is dissolved in water, they break down into their ions, which... 

Solutions of substances that are good conductors of electricity are called electrolytes. Sodium chloride, the major constituent of seawater, is a strong electrolyte. Most salts, as well as strong acids and bases, are strong electrolytes because they remain in solution primarily in ionic (charged) forms. Weak acids and bases are weak electrolytes because they tend to remain in nonionic forms. Pure water is a nonconductor of electricity. 

When an acid in solution is exactly neutralized with a base the resulting solution corresponds to a solution of the salt of the acid-base pair. This is a situation which frequently arises in analytical procedures and the calculation of the exact pH of such a solution may be of considerable importance. The neutralization point or end point in an acid-base titration is a particular example (Chapter 5). Salts may in all cases be regarded as strong electrolytes so that a salt AB derived from acid AH and base B will dissociate completely in solution. If the acid and base are strong, no further reaction is likely and the solution pH remains unaffected by the salt. However if either or both acid and base are weak a more complex situation will develop. It is convenient to consider three separate cases, (a) weak acid-strong base, (b) strong acid-weak base and (c) weak acid-weak base. 

At the microscopic level, the Arrhenius theory defines acids as substances which, when dissolved in water, yield the hydronium ion (H30+) or H+(aq). Bases are defined as substances which, when dissolved in water, yield the hydroxide ion (OH). Acids and bases may be strong (as in strong electrolytes), dissociating completely in water, or weak (as in weak electrolytes), partially dissociating in water. (We will see the more useful Brpnsted-Lowry definitions of acids and bases in Chapter 15.) Strong acids include ... 

In this chapter, you learned about solutions and how to use molarity to express the concentration of solutions. You also learned about electrolytes and nonelectrolytes. Using a set of solubility rules allows you to predict whether or not precipitation will occur if two solutions are mixed. You examined the properties of acids and bases and the neutralization reactions that occur between them. You then learned about redox reactions and how to use an activity table to predict redox reactions. You learned about writing net ionic equations. Finally, you learned how to use the technique of titrations to determine the concentration of an acid or base solution. 

B—A (nitrous acid) and D (acetic acid) are weak acids, and E (ammonia) is a weak base. Weak acids and bases are weak electrolytes. C (ethanol) is a nonelectrolyte. Potassium nitrate (B) is a water-soluble ionic compound. 

Strong acids and bases (and strong electrolytes) dissociate completely in water. Therefore, you can use the concentrations of these compounds to determine the concentrations of the ions they form in aqueous solutions. You cannot, however, use the concentrations of weak acids, bases, and electrolytes in the same way. Their solutions contain some particles that have not dissociated into ions. Nevertheless, important changes in [HsOT and [OH ] take place because dissolved ions affect the dissociation of water. 

Although these effects are often collectively referred to as salt effects, lUPAC regards that term as too restrictive. If the effect observed is due solely to the influence of ionic strength on the activity coefficients of reactants and transition states, then the effect is referred to as a primary kinetic electrolyte effect or a primary salt effect. If the observed effect arises from the influence of ionic strength on pre-equilibrium concentrations of ionic species prior to any rate-determining step, then the effect is termed a secondary kinetic electrolyte effect or a secondary salt effect. An example of such a phenomenon would be the influence of ionic strength on the dissociation of weak acids and bases. See Ionic Strength... 

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