Dissociation is the separation of a molecule into ions, and it is a key factor for evaluating the "strength" of acids and bases. The more a substance is prone to dissociation, the better it can conduct an electric current, because the separation of charges provides a "pathway" for the current's flow. A substance that dissociates completely, or almost completely, is called a strong electrolyte, whereas one that dissociates only slightly (or not at all) is designated as a weak electrolyte.
The terms "weak" and "strong" are also applied to acids and bases. For instance, vinegar is a weak acid, because it dissociates only slightly, and therefore conducts little electric current. By contrast, hydrochloric acid (HCl) is a strong acid, because it dissociates almost completely into positively charged hydrogen ions and negatively charged chlorine ones. Represented symbolically, this is: HCl →H + + Cl − .
It may seem a bit backward that a strong acid or base is one that "falls apart," while the weak one stays together. To understand the difference better, let us return to the reaction described earlier, in which an acid in aqueous solution reacts with water to produce a base in aqueous solution, along with hydronium: HA( aq ) + H 2 O( l ) →H 3 O + ( aq ) + A − ( aq ). Instead of using the generic symbols HA and A − , however, let us substitute hydrochloric acid (HCl) and chloride (Cl − ) respectively.
The reaction HCl( aq ) + H 2 O( l ) →H 3 O + ( aq ) + Cl − ( aq ) is a reversible one, and for that reason, the symbol for chemical equilibrium (⇋) can be inserted in place of the arrow pointing to the right. In other words, the substances on the right can just as easily react, producing the substances on the left. In this reverse reaction, the reactants of the forward reaction would become products, and the products of the forward reaction serve as the reactants.
However, the reaction described here is not perfectly reversible, and in fact the most proper chemical symbolism would show a longer arrow pointing to the right, with a shorter arrow pointing to the left. Due to the presence of a strong electrolyte, there is more forward "thrust" to this reaction.
Because it is a strong acid, the hydrogen chloride in solution is not a set of molecules, but a collection of H + and Cl − ions. In the reaction, the weak Cl − ions to the right side of the equilibrium symbol exert very little attraction for the H + ions. Instead of bonding with the chloride, these hydrogen ions join the water (a stronger base) to form hydronium.
The chloride, incidentally, is the conjugate base of the hydrochloric acid, and this illustrates another principal regarding the "strength" of electrolytes: a strong acid produces a relatively weak conjugate base. Likewise, a strong base produces a relatively weak conjugate acid.
There are only a few strong acids and bases, which are listed below:
Virtually all others are weak acids or bases, meaning that only a small percentage of molecules in these substances ionize by dissociation. The concentrations of the chemical species involved in the dissociation of weak acids and bases are mathematically governed by the equilibrium constant K..
Neutralization is the process whereby an acid and base react with one another to form a salt and water. The simplest example of this occurs in the reaction discussed earlier, in which hydrochloric acid or HCl( aq ) reacts with the base sodium hydroxide, designated as NaOH( aq ), in an aqueous solution. The result is sodium chloride, or common table salt (NaCl[ aq ]) and H 2 O. This equation is written thus: HCl( aq ) + NaOH( aq ) →NaCl( aq ) + H 2 O.
The human stomach produces hydrochloric acid, commonly known as "stomach acid." It is generated in the digestion process, but when a person eats something requiring the stomach to work overtime in digesting it—say, a pizza—the stomach may generate excess hydrochloric acid, and the result is "heartburn." When this happens, people often take antacids, which contain a base such as aluminum hydroxide (Al[OH] 3 ) or magnesium hydroxide (Mg[OH] 2 ).
When a person takes an antacid, the reaction leads to the creation of a salt, but not the salt with which most people are familiar—NaCl. As shown above, that particular salt is the product of a reaction between hydrochloric acid and sodium hydroxide, but a person who ingested sodium hydroxide (a substance used to unclog drains and clean ovens) would have much worse heartburn than before! In any case, the antacid reacts with the stomach acid to produce a salt, as well as water, and thus the acid is neutralized.
When land formerly used for mining is reclaimed, the acidic water in the area must be neutralized, and the use of calcium oxide (CaO) as a base is one means of doing so. Acidic soil, too, can be neutralized by the introduction of calcium carbonate (CaCO 3 ) or limestone, along with magnesium carbonate (MgCO 3 ). If soil is too basic, as for instance in areas where there has been too little precipitation, acid-like substances such as calcium sulfate or gypsum (SaSO 4 ) can be used. In either case, neutralization promotes plant growth.
One of the most important applications of neutralization is in titration, the use of a chemical reaction to determine the amount of a chemical substance in a sample of unknown purity. In a typical form of neutralization titration, a measured amount of an acid is added to a solution containing an unknown amount of a base. Once enough of the acid has been added to neutralize the base, it is possible to determine how much base exists in the solution.
Titration can also be used to measure pH ("power of hydrogen") level by using an acid-base indicator. The pH scale assigns values ranging from 0 (a virtually pure acid) to 14 (a virtually pure base), with 7 indicating a neutral substance. An acid-base indicator such as litmus paper changes color when it neutralizes the solution.
The transition interval (the pH at which the color of an indicator changes) is different for different types of indicators, and thus various indicators are used to measure substances in specific pH ranges. For instance, methyl red changes from red to yellow across a pH range of 4.4 to 6.2, so it is most useful for testing a substance suspected of being moderately acidic.
A buffered solution is one that resists a change in pH even when a strong acid or base is added to it. This buffering results from the presence of a weak acid and a strong conjugate base, and it can be very important to organisms whose cells can endure changes only within a limited range of pH values. Human blood, for instance, contains buffering systems, because it needs to be at pH levels between 7.35 and 7.45.
The carbonic acid-bicarbonate buffer system is used to control the pH of blood. The most important chemical equilibria (that is, reactions involving chemical equilibrium) for this system are: H + + HCO 3 − ⇋ H 2 CO 3 ⇋ H 2 0 + CO 2 . In other words, the hydrogen ion (H + ) reacts with the hydrogen carbonate ion (HCO 3 − ) to produce carbonic acid (H 2 CO 3 ). The latter is in equilibrium with the first set of reactants, as well as with water and carbon dioxide in the forward reaction.
The controls the pH level by changing the concentration of carbon dioxide by exhalation. In accordance with Le Châtelier's principle, this shifts the equilibrium to the right, consuming H + ions. In normal blood plasma, the concentration of HCO 3 − is about 20 times as great as that of H 2 CO 3 , and this large concentration of hydrogen carbonate gives the buffer a high capacity to neutralize additional acid. The buffer has a much lower capacity to neutralize bases because of the much smaller concentration of carbonic acid.
Water is an amphoteric substance; in other words, it can serve either as an acid or a base. When water experiences ionization, one water molecule serves as a Brønsted-Lowry acid, donating a proton to another water molecule—the Brønsted-Lowry base. This results in the production of a hydroxide ion and a hydronium ion: H 2 O( l ) + H 2 O( l ) ⇋ H 3 O + ( aq ) + OH − ( aq ).
This equilibrium equation is actually one in which the tendency toward the reverse reaction is much greater; therefore the equilibrium symbol, if rendered in its most proper form, would show a much shorter arrow pointing toward the right. In water purified by distillation, the concentrations of hydronium (H 3 O + ) and hydroxide (OH − ) ions are equal. When multiplied by one another, these yield the constant figure 1.0 · 10 −14 , which is the equilibrium constant for water. In fact, this constant—denoted as K w —is called the ion-product constant for water.
Because the product of these two concentrations is always the same, this means that if one of them goes up, the other one must go down in order to yield the same constant. This explains the fact, noted earlier, that water can serve either as an acid or base—or, if the concentrations of hydronium and hydroxide ions are equal—as a neutral substance. In situations where the concentration of hydronium is higher, and the hydroxide concentration automatically decreases, water serves as an acid. Conversely, when the hydroxide concentration is high, the hydronium concentration decreases correspondingly, and the water is a base.
"Acids and Bases." Vision Learning (Web site). <http://www.visionlearning.com/library/science/chemistry-2/CHE2.2-acid_base.htm> (June 7, 2001).
"Acids, Bases, and Chemical Reactions" Open Access College (Web site). <http://oac.schools.sa.edu.au/8-10science/acids.htm> (June 7, 2001).
"Acids, Bases, pH." About.com (Web site). <http://chemistry.about.com/science/chemistry/cs/acidsbasesph/ (June 7, 2001).
"Acids, Bases, and Salts" (Web site). <http://edie.cprost.sfu.ca/~rhlogan/ionic_eq.html> (June 7, 2001).
"Junior Part: Acids, Bases, and Salts" (Web site). <http://www.rjclarkson.demon.co.uk/junior/junior4.htm> (June 7, 2001).
Knapp, Brian J. Acids, Bases, and Salts. Danbury, CT: Grolier Educational, 1998.
Moje, Steven W. Cool Chemistry: Great Experiments with Simple Stuff. New York: Sterling Publishing Company, 1999.
Walters, Derek. Chemistry. Illustrated by Denis Bishop and Jim Robins. New York: F. Watts, 1982.
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.