16. Ions and Equilibrium;
       Acids and Bases
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       Measuring pH: Acid-Base Indicators

An acid-base indicator is a weak acid (or a weak base) that has different colors in its un-ionized and ionized states. Most indicators are aromatic molecules that have delocalized electrons, and in Chapter 9 we saw the reason for their color changes. The equilibrium

with HIn representing the acid form of the indicator compound, is shifted to the left by an excess of acid, and to the right by an excess of base. The ratio of basic to acidic form of the indicator is linked to the pH by the now familiar expression

in which pKa is the acid-dissociation constant for the weak indicator acid, HIn. The eye is sensitive to color changes over approximately a 1:10 to 10:1 concentration ratio, meaning that visible color changes in an indicator occur in a pH range of around 2 units, centered on the indicator's own pKa. Litmus paper changes from red in acid to blue in base, in the pH range 5-8. Phenolphthalein solution added in minute quantities to the solution being tested or titrated changes from colorless (acid) to red (base) in the range of pH 8-10 because it has a pKa around 9. Other common indicators, color changes, and useful pH ranges are shown in the diagram on page 30.



The most familiar salt, sodium chloride, is so soluble that we tend to think that all salts are equally soluble. This is far from true, and many salts are quite insoluble in water.

Solubility is the result of competition between the mutual attractions of ions in the crystal, and hydration of individual ions by solvent molecules. Both of these processes involve large energies, and solubility depends on the frequently quite small difference between them. It is difficult to calculate crystal-lattice energies and hydration energies accurately enough to predict whether the difference between them will be positive or negative. Although we understand the forces at work when a salt dissolves, it is not easy to predict whether a given salt will dissolve or not.

There are a few common-sense principles that help in a general way. A crystal is held together by electrostatic forces between oppositely charged ions. Crystals with small ions that can be packed close together generally are harder to pull apart than crystals with large ions.

Hence fluorides (F-) and hydroxides (OH-) tend to be less soluble than nitrates (NO3-) and perchlorates (ClO4-) with the same positive ion; chlorides (Cl-) are intermediate in solubility.

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