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 negativelycharged 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:

Chemically, acids and bases may be considered 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 where 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 abilities to change colors of certain other chemicals. Historically, 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. Such materials are called indicators.

Classic definition of acids and bases

Although acids and bases have been known since prehistoric times (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. H2SO4 is an example. Triprotic acids, like H3PO4, can release three hydrogen ions into solution. Acetic acid has the formula HC2H3O2 and is a monoprotic acid because it is composed of one H+ion and one acetate ion, C2H3O2-. 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 act 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 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. There are only a few strong acids: hydrochloric acid (HCl), hydrobromic acid (HBr), hydriodic acid (HI), perchloric acid (HClO4), nitric acid (HNO3), and sulfuric acid (H2SO4). Similarly, there are only a few strong bases: lithium hydroxide (LiOH), sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca[OH]2), strontium hydroxide (Sr[OH]2), and barium hydroxide (Ba[OH]2).

These strong acids and bases are 100% ionized in aqueous solution. All other Arrhenius acids and bases are weak acids and bases. For example, acetic acid (HC2H3O2) and oxalic acid (H2C2O4) are weak acids, while iron hydroxide, Fe(OH)3, and ammonium hydroxide, NH4OH (which is actually just ammonia , NH3, dissolved in water), are examples of weak bases. 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 unionized 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 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 solution. In many other cases, water is indeed the solvent. In many cases, however, water is not the solvent. What was necessary was to formulate a definition of acid and base that were 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) and is called the Brønsted-Lowry definition of acids and bases. (Bjerrum seems to have been forgotten.) 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.

The Brønsted-Lowry definition includes all Arrhenius acids and bases, since the hydrogen ion is a proton donor (in fact, it is a proton) and a hydroxide ion accepts a proton to form water:

But the Brønsted-Lowry definition also includes chemical species that are not Arrhenius-type acids or bases. The classic example is ammonia, NH3. Ammonia dissolves in water to make a slightly basic solution even though ammonia does not contain OH– ions. What is happening is that an ammonia molecule is accepting a proton from a water molecule to make an ammonium ion (NH4+) and a hydroxide ion:

In essence, the water molecule is donating a proton to the ammonia molecule. The water molecule is therefore acting as the Brønsted-Lowry acid and the ammonia molecule is acting as the Brønsted-Lowry base.

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 H3O+. 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.

Acid	Name	Use	Base	Name	Use
HCl	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

Chemical reactions can go forward or backward; when the rates of the reverse reactions are equal, it is at chemical equilibrium. It can be shown that each side of the equilibrium has a Brønsted-Lowry acid and base. For example:

On each side of the reaction there is an acid and a base. The NH +4 ion is an acid because in the reverse reaction it donates a proton (H+) to the OH-ion to form NH3 and H2O. With respect to the reaction above, the H2O and OH- species make up an acid-base pair, called a conjugate acid-base pair, while the NH3 and NH4+ species make up another conjugate acid-base pair. All Brønsted-Lowry acid-base reactions can be separated into reactions between two conjugate acid-base pairs. The conjugate acid always has one more H+ than the conjugate base.

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 reacts with another compound is what determines whether it is an acid or a base.

To show that the Lewis definition is not in conflict with previous definitions of acid and base, consider the fundamental acid-base reaction of H+ with OH- to give H2O. The oxygen atom in the hydroxide ion has three unbonded electron pairs around it, and during the course of the reaction one of those electron pairs is "donated" to the hydrogen ion, making a chemical bond . Thus, OH-is the electron pair donor and the Lewis base, whereas H+ is the electron pair acceptor and, therefore, the Lewis acid. These assignments are consistent with both the Arrhenius definition and the Brønsted-Lowry definitions of acid and base.

However, the Lewis acid/base definition is much broader than the previous two definitions. Consider the reaction of BF3 and NH3 in the gas phase, in which NH3 is donating an electron pair to the BF3molecule:

Compounds like F3BNH3 are stable and can be purchased as solutions in organic solvents or even as pure compounds. In the above chemical reaction, BF3 is accepting an electron pair and therefore is the Lewis acid; NH3 is donating the electron pair and so is the Lewis base. However, in this case neither the Arrhenius definition nor the Brønsted-Lowry definition are applicable. Therefore, while the Lewis acid/base definition includes acids and bases from the other two definitions, it expands the definitions to include compounds that are not otherwise considered "classic" acids and bases.

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 of certain groupings of atoms within an organic molecule called functional groups as acidic or basic. 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 and is produced by some ants and causes their bites to sting. Another example is acetic acid, CH3COOH, 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. Their utility is illustrated by the fact that three of the top 10 chemicals produced in the US in 1994 are acids or bases: sulfuric acid (#1, 89 billion lbs/40 billion kg produced), sodium hydroxide (#8, 26 billion lbs/12 billion kg produced), and phosphoric acid (#9, 25 billion lbs/11 billion kg produced). Weak acids and bases have specific uses in society which are so variable that the specific compound entry should be consulted.

See also Acetic acid; Alkaloid; Carboxylic acids; Citric acid; Neutralization; Nitric acid; Sodium hydroxide; Sulfuric acid.

Resources

books

Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002.

Scorpio, Ralph. Fundamental of Acids, Bases, Buffers & TheirApplication to Biochemical Systems. Falls Church, VA: Kendall/Hunt, 2000.

Snyder, C.H. The Extraordinary Chemistry of Ordinary Things. 4th ed. New York: John Wiley and Sons, 2002.

David W. Ball

KEY TERMS

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Functional group

—In organic chemistry, certain 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.