Acids and bases
Acids and bases
Classic definition of acids and bases
Strong and weak acids and bases
Brønsted-Lowry definition of acids and bases
Lewis definition of acids and bases
Acids and bases are chemical compounds that have certain specific properties in aqueous solutions. In most chemical circumstances, acids are chemicals that produce positively-charged hydrogen ions (H+) in water, while bases are chemicals that produce negatively-charged hydroxide ions (OH–) in water. Bases are sometimes called alkalis. Acids and bases react with each other in a reaction called neutralization. In a neutralization reaction, the hydrogen ion and the hydroxide ion react to form a molecule of water:
H++OH–→H2O
Chemically, acids and bases are considered to be opposites of each other. The concept of acids and bases is so important in chemistry that there are several useful definitions of “acid” and “base” that pertain to different chemical environments, although the definition above is the most common one.
Acids and bases have some general properties. Many acids have a sour taste. Citric acid, found in oranges and lemons, is one example in which the sour taste is related to the fact that the chemical is an acid. Molecules that are bases usually have a bitter taste, like caffeine. Bases make solutions that are slippery. Many acids will react with metals to dissolve the metal and, at the same time, generate hydrogen gas, H2. Perhaps the most obvious behavior of acids and bases is their ability to change the colors of certain other chemicals. Historically, to determine if a solution is acidic or basic, an extract of lichens (V. lecanora and V. rocella ) called litmus has been used, since it turns blue in the presence of bases and red in the presence of acids. Litmus paper is still commonly used to indicate whether a compound is an acid or a base. Extracts made from red onions, red cabbage, and many other fruits and vegetables change colors in the presence of acids and bases. These materials are called indicators.
Classic definition of acids and bases
Although acids and bases have been known for centuries (vinegar, for example, is an acid), the first attempt to define what makes a compound an acid or a base was made by the Swedish chemist Svante Arrhenius (1859-1927), who proposed the definition that an acid was any compound that produced hydrogen ions, H+, when dissolved in water, and a base was any compound that produced hydroxide ions, OH-, when dissolved in water. Although this was (and still is) a very useful definition, it has two major limitations. First, it was limited to water, or aqueous, solutions. Second, it practically limited acids and bases to ionic compounds that contained the H+ ion or the OH- ion (compounds like hydrochloric acid, HCl, or sodium hydroxide, NaOH). Limited though it might be, it was an important step in the understanding of chemistry in solutions, and for his work on solution chemistry Arrhenius was awarded the 1903 Nobel Prize in chemistry.
Many common acids and bases are consistent with the Arrhenius definition. The following table shows a few common acids and bases and their uses. In all cases it is assumed that the acid or base is dissolved in water.
Many acids release only a single hydrogen ion per molecule into solution. Such acids are called monoprotic. Examples include hydrochloric acid, HCl, and nitric acid, HNO3. Diprotic acids can release two hydrogen ions per molecule. H2 SO4 is an example. Triprotic acids, like H3 PO4, can release three hydrogen ions into solution. Acetic acid has the formula HC2 H3 O2 and is a monoprotic acid because it is composed of one H+ ion and one acetate ion, C2 H3 O2–. The three hydrogen atoms in the acetate ion do not act as acids.
Strong and weak acids and bases
An important consideration when dealing with acids and bases is their strength; that is, how chemically reactive they are as acids and bases. The strength of an acid or base is determined by the degree of ionization of the acid or base in solution—that is, the percentage of dissolved acid or base molecules that
| Acids and Bases and Their Common Uses. (Thomson Gale.) | |||||
|---|---|---|---|---|---|
| Acids and bases | |||||
| Acid | Name | Use | Base | Name | Use |
| HCI | hydrochloric acid | cleaning, drugs, plastics | NaOH | sodium hydroxide | drain cleaner, soap |
| H2SO4 | sulfuric acid | chemical synthesis, batteries | KOH | potassium hydroxide | soaps |
| HC2H3O2 | acetic acid | vinegar | Mg(OH)2 | magnesium hydroxide | antacids |
release hydrogen or hydroxide ions. If all of the dissolved acid or base separates into ions, it is called a strong acid or strong base. Otherwise, it is a weak acid or weak base.
Strong acids and bases are 100% ionized in aqueous solution; in other words they completely dissociate to form ions. In contrast, weak acids or bases do not completely ionize. The percentage of the acid and base molecules that are ionized in solution varies and depends on the concentration of the acid. For example, a 2% solution of acetic acid in water, which is about the concentration found in vinegar, is only 0.7% ionized. This means that fully 99.3% of the acetic acid molecules are not ionized and exist in solution as the complete acetic acid molecule.
Brønsted-Lowry definition of acids and bases
Although the Arrhenius definitions of acids and bases are the simplest and most useful, they are not the most widely applicable. Some compounds, like ammonia, NH3, act like bases in aqueous solution even though they are not hydroxide-containing compounds. Also, the Arrhenius definition assumes that the acid-base reactions are occurring in aqueous (water-based) solutions. In some cases, water is indeed the solvent, yet in other cases, water is not the solvent; rather an alcohol or some other liquid is involved. Thus, it was necessary to formulate definitions of acids and bases independent of the solvent and the presence of H+ and OH–ions.
Such a definition was proposed in 1923 by English chemist Thomas Lowry (1874-1936) and Danish chemists J. N. Brønsted (1879-1947) and N. Bjerrum (1879-1958); it is called the Brønsted-Lowry definition of acids and bases. The central chemical species of this definition is H+, which consists merely of a proton. By the Brønsted-Lowry definition, an acid is any chemical species that donates a proton to another chemical species. Conversely, a base is any chemical species that accepts a proton from another chemical species. Simply put, a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor.
In order to better understand the Brønsted-Lowry definition, it needs to be understood what is meant by a proton. The descriptions proton donor and proton acceptor are easy to remember, but are there actually bare protons floating around in solution? Not really. In aqueous solution, the protons are attached to the oxygen atoms of water molecules, giving them a positive charge. This species is called the hydronium ion and has the chemical formula H3 O+. It is more accurate to use the hydronium ion instead of the bare hydrogen ion when writing equations for chemical reactions between acids and bases in aqueous solution. For example, the reaction between the hydronium ion and the hydroxide ion, the typical Arrhenius acid-base reaction, would produce two molecules of water.
Lewis definition of acids and bases
The Brønsted-Lowry acid-base definition, while broader than the Arrhenius definition, is still limited to hydrogen-containing compounds, and is dependent on a hydrogen ion (that is, a proton) transferring from one molecule to another. Ultimately, a definition of acid and base that is completely independent of the presence of a hydrogen atom is necessary.
Such a definition was provided in 1923 by American chemist Gilbert N. Lewis (1875-1946). Instead of focusing on protons, Lewis’s definition focuses on electron pairs. Since all compounds contain electron pairs, the Lewis definition is applicable to a wide range of chemical reactions.
A Lewis acid is defined as the reactant in a chemical reaction that accepts an electron pair from another reactant. A Lewis base is defined as the reactant in a chemical reaction that donates an electron pair to another reactant. Like the Brønsted-Lowry definition of acids and bases, the Lewis definition is reaction-dependent. A compound is not an acid or base in its own right; rather, how that compound
KEY TERMS
Functional group —Specific groupings of atoms in a molecule.
Ionic compound —A compound consisting of positive ions (usually metal ions) and negative ions (nonmetal ions) held together by electrostatic attraction.
reacts with another compound is what determines whether it is an acid or a base.
Organic acids
Organic chemistry is the study of compounds of the element carbon. Organic chemistry uses the ideas of acids and bases in two ways. The more general way is that the concept of Lewis acids and bases is used to classify organic chemical reactions as acid/base reactions because the donation of electron pairs is quite common.
The second way that organic chemistry uses the concepts of acids and bases is in the definition, as acidic or basic, of certain groupings (called functional groups) of atoms within an organic molecule. An organic base is, in the true Lewis base style, any molecule with electron pairs that can be donated. The most common organic base involves a nitrogen atom, N, bonded to carbon-containing groups. One important class of such compounds is known as amines. In these compounds, the nitrogen atom has an unbonded electron pair that it can donate as it reacts as a Lewis base. Several of these compounds are gases and have a somewhat putrid, fish-like odor. These compounds are relatively simple molecules. There are larger organic molecules, including many of natural origin, that contain a nitrogen atom and so have certain base-like properties. These compounds are called alkaloids. Examples include quinine, caffeine, strychnine, nicotine, morphine, and cocaine.
Organic chemistry uses the acid concept not only in the definition of the Lewis acid but also by defining a particular collection of atoms as an acid functional group. Any organic molecule containing a carboxyl group, –COOH, is called a carboxylic acid. (Non-organic acids are sometimes called mineral acids). Examples include formic acid, which has the formula HCOOH. This acid is produced by some ants and causes their bites to sting. Another example is acetic acid, CH3 COOH, which is the acid in vinegar.
Uses of acids and bases
Many specific uses of acids and bases have been discussed above. Generally, strong acids and bases are used for cleaning and, most importantly, for synthesizing other compounds.
Resources
BOOKS
Gonick, Larry, and Craig Criddle. The Cartoon Guide to Chemistry. New York: Collins, 2005.
Moore, John T. Chemistry for Dummies. New York: For Dummies, 2002.
Preston, Richard A. Acid-Base, Fluids, and Electrolytes Made Ridiculously Simple. Miami: MedMaster, 2002.
David W. Ball