ACIDS AND BASES

CONCEPT

The name "acid" calls to mind vivid sensory images—of tartness, for instance, if the acid in question is meant for human consumption, as with the citric acid in lemons. On the other hand, the thought of laboratory-and industrial-strength substances with scary-sounding names, such as sulfuric acid or hydrofluoric acid, carries with it other ideas—of acids that are capable of destroying materials, including human flesh. The name "base," by contrast, is not widely known in its chemical sense, and even when the older term of "alkali" is used, the sense-impressions produced by the word tend not to be as vivid as those generated by the thought of "acid." In their industrial applications, bases too can be highly powerful. As with acids, they have many household uses, in substances such as baking soda or oven cleaners. From a taste standpoint, (as anyone who has ever brushed his or her teeth with baking soda knows), bases are bitter rather than sour. How do we know when something is an acid or a base? Acid-base indicators, such as litmus paper and other materials for testing pH, offer a means of judging these qualities in various substances. However, there are larger structural definitions of the two concepts, which evolved in three stages during the late nineteenth and early twentieth centuries, that provide a more solid theoretical underpinning to the understanding of acids and bases.

HOW IT WORKS

Introduction to Acids and Bases

Prior to the development of atomic and molecular theory in the nineteenth century, followed by the discovery of subatomic structures in the late nineteenth and early twentieth centuries, chemists could not do much more than make measurements and observations. Their definitions of substances were purely phenomenological—that is, the result of experimentation and the collection of data. From these observations, they could form general rules, but they lacked any means of "seeing" into the atomic and molecular structures of the chemical world.

The phenomenological distinctions between acids and bases, gathered by scientists from ancient times onward, worked well enough for many centuries. The word "acid" comes from the Latin term acidus, or "sour," and from an early period, scientists understood that substances such as vinegar and lemon juice shared a common acidic quality. Eventually, the phenomenological definition of acids became relatively sophisticated, encompassing such details as the fact that acids produce characteristic colors in certain vegetable dyes, such as those used in making litmus paper. In addition, chemists realized that acids dissolve some metals, releasing hydrogen in the process.

WHY "BASE" AND NOT "ALKALI"?

The word "alkali" comes from the Arabic al-qili, which refers to the ashes of the seawort plant. The latter, which typically grows in marshy areas, was often burned to produce soda ash, used in making soap. In contrast to acids, bases—caffeine, for example—have a bitter taste, and many of them feel slippery to the touch. They also produce characteristic colors in the vegetable dyes of litmus paper, and can be used to promote certain chemical reactions. Note that today chemists use the word "base" instead of "alkali," the reason being that the latter term has a narrower meaning: all alkalies are bases, but not all bases are alkalies.

Originally, "alkali" referred only to the ashes of burned plants, such as seawort, that contained either sodium or potassium, and from which the oxides of sodium and potassium could be obtained. Eventually, alkali came to mean the soluble hydroxides of the alkali and alkaline earth metals. This includes sodium hydroxide, the active ingredient in drain and oven cleaners; magnesium hydroxide, used for instance in milk of magnesia; potassium hydroxide, found in soaps and other substances; and other compounds. Broad as this range of substances is, it fails to encompass the wide array of materials known today as bases—compounds which react with acids to form salts and water.

Toward a Structural Definition

The reaction to form salts and water is, in fact, one of the ways that acids and bases can be defined. In an aqueous solution, hydrochloric acid and sodium hydroxide react to form sodium chloride—which, though it is suspended in an aqueous solution, is still common table salt—along with water. The equation for this reaction is HCl(aq ) + NaOH(aq ) →H₂O + NaCl(aq ). In other words, the sodium (Na) ion in sodium hydroxide switches places with the hydrogen ion in hydrochloric acid, resulting in the creation of NaCl (salt) along with water.

But why does this happen? Useful as this definition regarding the formation of salts and water is, it is still not structural—in other words, it does not delve into the molecular structure and behavior of acids and bases. Credit for the first truly structural definition of the difference goes to the Swedish chemist Svante Arrhenius (1859-1927). It was Arrhenius who, in his doctoral dissertation in 1884, introduced the concept of an ion, an atom possessing an electric charge.

His understanding was particularly impressive in light of the fact that it was 13 more years before the discovery of the electron, the subatomic particle responsible for the creation of ions. Atoms have a neutral charge, but when an electron or electrons depart, the atom becomes a positive ion or cation. Similarly, when an electron or electrons join a previously uncharged atom, the result is a negative ion or anion. Not only did the concept of ions greatly influence the future of chemistry, but it also provided Arrhenius with the key necessary to formulate his distinction between acids and bases.

The Arrhenius Definition

Arrhenius observed that molecules of certain compounds break into charged particles when placed in liquid. This led him to the Arrhenius acid-base theory, which defines an acid as any compound that produces hydrogen ions (H⁺) when dissolved in water, and a base as any compound that produces hydroxide ions (OH⁻) when dissolved in water.

This was a good start, but two aspects of Arrhenius's theory suggested the need for a definition that encompassed more substances. First of all, his theory was limited to reactions in aqueous solutions. Though many acid-base reactions do occur when water is the solvent, this is not always the case.

Second, the Arrhenius definition effectively limited acids and bases only to those ionic compounds, such as hydrochloric acid or sodium hydroxide, which produced either hydrogen or hydroxide ions. However, ammonia, or NH₃, acts like a base in aqueous solutions, even though it does not produce the hydroxide ion. The same is true of other substances, which behave like acids or bases without conforming to the Arrhenius definition.

These shortcomings pointed to the need for a more comprehensive theory, which arrived with the formulation of the Brønsted-Lowry definition by English chemist Thomas Lowry (1874-1936) and Danish chemist J. N. Brønsted (1879-1947). Nonetheless, Arrhenius's theory represented an important first step, and in 1903, he was awarded the Nobel Prize in Chemistry for his work on the dissociation of molecules into ions.

The BrØnsted-Lowry Definition

The Brønsted-Lowry acid-base theory defines an acid as a proton (H⁺) donor, and a base as a proton acceptor, in a chemical reaction. Protons are represented by the symbol H⁺, and in representing acids and bases, the symbols HA and A⁻, respectively, are used. These symbols indicate that an acid has a proton it is ready to give away, while a base, with its negative charge, is ready to receive the positively charged proton.

Though it is used here to represent a proton, it should be pointed out that H⁺ is also the hydrogen ion—a hydrogen atom that has lost its sole electron and thus acquired a positive charge. It is thus really nothing more than a lone proton, but this is the one and only case in which an atom and a proton are exactly the same thing. In an acid-base reaction, a molecule of acid is "donating" a proton, in the form of a hydrogen ion. This should not be confused with a far more complex process, nuclear fusion, in which an atom gives up a proton to another atom.

AN ACID-BASE REACTION IN BRØNSTED-LOWRY THEORY.

The most fundamental type of acid-base reaction in Brønsted-Lowry theory can be symbolized thus HA(aq ) + H₂O(l ) →H₃O⁺(aq ) + A⁻(aq ). The first acid shown—which, like three of the four "players" in this equation, is dissolved in an aqueous solution—combines with water, which can serve as either an acid or a base. In the present context, it functions as a base.

Water molecules are polar, meaning that the negative charges tend to congregate on one end of the molecule with the oxygen atom, while the positive charges remain on the other end with the hydrogen atoms. The Brønsted-Lowry model emphasizes the role played by water, which pulls the proton from the acid, resulting in the creation of H₃O⁺, known as the hydronium ion.

The hydronium ion produced here is an example of a conjugate acid, an acid formed when a base accepts a proton. At the same time, the acid has lost its proton, becoming A⁻, a conjugate base—that is, the base formed when an acid releases a proton. These two products of the reaction are called a conjugate acid-base pair, a term that refers to two substances related to one another by the donating of a proton.

Brønsted and Lowry's definition represents an improvement over that of Arrhenius, because it includes all Arrhenius acids and bases, as well as other chemical species not encompassed in Arrhenius theory. An example, mentioned earlier, is ammonia. Though it does not produce OH⁻ ions, ammonia does accept a proton from a water molecule, and the reaction between these two (with water this time serving the function of acid) produces the conjugate acid-base pair of NH₄⁺ (an ammonium ion) and OH⁻. Note that the latter, the hydroxide ion, was not produced by ammonia, but is the conjugate base that resulted when the water molecule lost its H⁺ atom or proton.

The Lewis Definition

Despite the progress offered to chemists by the Brønsted-Lowry model, it was still limited to describing compounds that contain hydrogen. As American chemist Gilbert N. Lewis (1875-1946) recognized, this did not encompass the full range of acids and bases; what was needed, instead, was a definition that did not involve the presence of a hydrogen atom.

Lewis is particularly noted for his work in the realm of chemical bonding. The bonding of atoms is the result of activity on the part of the valence electrons, or the electrons at the "outside" of the atom. Electrons are arranged in different ways, depending on the type of bonding, but they always bond in pairs.

According to the Lewis acid-base theory, an acid is the reactant that accepts an electron pair from another reactant in a chemical reaction, while a base is the reactant that donates an electron pair to another reactant. As with the Brønsted-Lowry definition, the Lewis definition is reaction-dependant, and does not define a compound as an acid or base in its own right. Instead, the manner in which the compound reacts with another serves to identify it as an acid or base.

AN IMPROVEMENT OVER ITS PREDECESSORS.

The beauty of the Lewis definition lies in the fact that it encompasses all the situations covered by the others—and more. Just as Brønsted-Lowry did not disprove Arrhenius, but rather offered a definition that covered more substances, Lewis expanded the range of substances beyond those covered by Brønsted-Lowry. In particular, Lewis theory can be used to differentiate the acid and base in bond-producing chemical reactions where ions are not produced, and in which there is no proton donor or acceptor. Thus it represents an improvement over Arrhenius and Brønsted-Lowry respectively.

An example is the reaction of boron trifluoride (BF₃) with ammonia (NH₃), both in the gas phases, to produce boron trifluoride ammonia complex (F₃BNH₃). In this reaction, boron trifluoride accepts an electron pair and is therefore a Lewis acid, while ammonia donates the electron pair and is thus a Lewis base. Though hydrogen is involved in this particular reaction, Lewis theory also addresses reactions involving no hydrogen.

REAL-LIFE APPLICATIONS

pHand Acid-Base Indicators

Though chemists apply the sophisticated structural definitions for acids and bases that we have discussed, there are also more "hands-on" methods for identifying a particular substance (including complex mixtures) as an acid or base. Many of these make use of the pH scale, developed by Danish chemist SØren SØrensen (1868-1939) in 1909.

The term pH stands for "potential of hydrogen," and the pH scale is a means of determining the acidity or alkalinity of a substance. (Though, as noted, the term "alkali" has been replaced by "base," alkalinity is still used as an adjectival term to indicate the degree to which a substance displays the properties of a base.) There are theoretically no limits to the range of the pH scale, but figures for acidity and alkalinity are usually given with numerical values between 0 and 14.

THE MEANING OF pH VALUES.

A rating of 0 on the pH scale indicates a substance that is virtually pure acid, while a 14 rating represents a nearly pure base. A rating of 7 indicates a neutral substance. The pH scale is logarithmic, or exponential, meaning that the numbers represent exponents, and thus an increased value of 1 represents not a simple arithmetic addition of 1, but an increase of 1 power. This, however, needs a little further explanation.

The pH scale is actually based on negative logarithms for the values of H₃O⁺ (the hydronium ion) or H⁺ (protons) in a given substance. The formula is thus pH = −log[H₃O⁺] or −log[H⁺], and the presence of hydronium ions or protons is measured according to their concentration of moles per liter of solution.

pH VALUES OF VARIOUS SUBSTANCES.

The pH of a virtually pure acid, such as the sulfuric acid in car batteries, is 0, and this represents 1 mole (mol) of hydronium per liter (l) of solution. Lemon juice has a pH of 2, equal to 10⁻² mol/l. Note that the pH value of 2 translates to an exponent of −2, which, in this case, results in a figure of 0.01 mol/l.

Distilled water, a neutral substance with a pH of 7, has a hydronium equivalent of 10⁻⁷ mol/l. It is interesting to observe that most of the fluids in the human body have pH values in the neutral range blood (venous, 7.35; arterial, 7.45); urine (6.0—note the higher presence of acid); and saliva (6.0 to 7.4).

At the alkaline end of the scale is borax, with a pH of 9, while household ammonia has a pH value of 11, or 10⁻¹¹ mol/l. Sodium hydroxide, or lye, an extremely alkaline chemical with a pH of 14, has a value equal to 10⁻¹⁴ moles of hydronium per liter of solution.

LITMUS PAPER AND OTHER INDICATORS.

The most precise pH measurements are made with electronic pH meters, which can provide figures accurate to 0.001 pH. However, simpler materials are also used. Best known among these is litmus paper (made from an extract of two lichen species), which turns blue in the presence of bases and red in the presence of acids. The term "litmus test" has become part of everyday language, referring to a make-or-break issue—for example, "views on abortion rights became a litmus test for Supreme Court nominees."

Litmus is just one of many materials used for making pH paper, but in each case, the change of color is the result of the neutralization of the substance on the paper. For instance, paper coated with phenolphthalein changes from colorless to pink in a pH range from 8.2 to 10, so it is useful for testing materials believed to be moderately alkaline. Extracts from various fruits and vegetables, including red cabbages, red onions, and others, are also applied as indicators.

Some Common Acids and Bases

The tables below list a few well-known acids and bases, along with their formulas and a few applications

Common Acids

• Acetic acid (CH₃COOH): vinegar, acetate
• Acetylsalicylic acid (HOOCC₆H₄OOCCH₃): aspirin
• Ascorbic acid (H₂C₆H₆O₆): vitamin C
• Carbonic acid (H₂CO₃): soft drinks, seltzer water
• Citric acid (C₆H₈O₇): citrus fruits, artificial flavorings
• Hydrochloric acid (HCl): stomach acid
• Nitric acid (HNO₃): fertilizer, explosives
• Sulfuric acid (H₂SO₄): car batteries

Common Bases

• Aluminum hydroxide (Al[OH]₃): antacids, deodorants
• Ammonium hydroxide (NH₄OH): glass cleaner
• Calcium hydroxide (Ca[OH]₂): caustic lime, mortar, plaster
• Magnesium hydroxide (Mg[OH]₂): laxatives, antacids
• Sodium bicarbonate/sodium hydrogen carbonate (NaHCO₃): baking soda
• Sodium carbonate (Na₂CO₃): dish detergent
• Sodium hydroxide (NaOH): lye, oven and drain cleaner
• Sodium hypochlorite (NaClO): bleach

Of course these represent only a few of the many acids and bases that exist. Selected substances listed above are discussed briefly below.

Acids

ACIDS IN THE HUMAN BODY AND FOODS.

As its name suggests, citric acid is found in citrus fruits—particularly lemons, limes, and grapefruits. It is also used as a flavoring agent, preservative, and cleaning agent. Produced commercially from the fermentation of sugar by several species of mold, citric acid creates a taste that is both tart and sweet. The tartness, of course, is a function of its acidity, or a manifestation of the fact that it produces hydrogen ions. The sweetness is a more complex biochemical issue relating to the ways that citric acid molecules fit into the tongue's "sweet" receptors.

Citric acid plays a role in one famous stomach remedy, or antacid. This in itself is interesting, since antacids are more generally associated with alkaline substances, used for their ability to neutralize stomach acid. The fizz in Alka-Seltzer, however, comes from the reaction of citric acids (which also provide a more pleasant taste) with sodium bicarbonate or baking soda, a base. This reaction produces carbon dioxide gas. As a preservative, citric acid prevents metal ions from reacting with, and thus hastening the degradation of, fats in foods. It is also used in the production of hair rinses and low-pH shampoos and toothpastes.

The carboxylic acid family of hydrocarbon derivatives includes a wide array of substances—not only citric acids, but amino acids. Amino acids combine to make up proteins, one of the principal components in human muscles, skin, and hair. Carboxylic acids are also applied industrially, particularly in the use of fatty acids for making soaps, detergents, and shampoos.

SULFURIC ACID.

There are plenty of acids found in the human body, including hydrochloric acid or stomach acid—which, in large quantities, causes indigestion, and the need for neutralization with a base. Nature also produces acids that are toxic to humans, such as sulfuric acid.

Though direct exposure to sulfuric acid is extremely dangerous, the substance has numerous applications. Not only is it used in car batteries, but sulfuric acid is also a significant component in the production of fertilizers. On the other hand, sulfuric acid is damaging to the environment when it appears in the form of acid rain. Among the impurities in coal is sulfur, and this results in the production of sulfur dioxide and sulfur trioxide when the coal is burned. Sulfur trioxide reacts with water in the air, creating sulfuric acid and thus acid rain, which can endanger plant and animal life, as well as corrode metals and building materials.

Bases

The alkali metal and alkaline earth metal families of elements are, as their name suggests, bases. A number of substances created by the reaction of these metals with nonmetallic elements are taken internally for the purpose of settling gastric trouble or clearing intestinal blockage. For instance, there is magnesium sulfate, better known as Epsom salts, which provide a powerful laxative also used for ridding the body of poisons.

Aluminum hydroxide is an interesting base, because it has a wide number of applications, including its use in antacids. As such, it reacts with and neutralizes stomach acid, and for that reason is found in commercial antacids such as Di-Gel™, Gelusil™, and Maalox™. Aluminum hydroxide is also used in water purification, in dyeing garments, and in the production of certain kinds of glass. A close relative, aluminum hydroxychloride or Al₂(OH)₅Cl, appears in many commercial antiperspirants, and helps to close pores, thus stopping the flow of perspiration.

SODIUM HYDROGEN CARBONATE (BAKING SODA).

Baking soda, known by chemists both as sodium bicarbonate and sodium hydrogen carbonate, is another example of a base with multiple purposes. As noted earlier, it is used in Alka-Seltzer™, with the addition of citric acid to improve the flavor; in fact, baking soda alone can perform the function of an antacid, but the taste is rather unpleasant.

Baking soda is also used in fighting fires, because at high temperatures it turns into carbon dioxide, which smothers flames by obstructing the flow of oxygen to the fire. Of course, baking soda is also used in baking, when it is combined with a weak acid to make baking powder. The reaction of the acid and the baking soda produces carbon dioxide, which causes dough and batters to rise. In a refrigerator or cabinet, baking soda can absorb unpleasant odors, and additionally, it can be applied as a cleaning product.

SODIUM HYDROXIDE (LYE).

Another base used for cleaning is sodium hydroxide, known commonly as lye or caustic soda. Unlike baking soda, however, it is not to be taken internally, because it is highly damaging to human tissue—particularly the eyes. Lye appears in drain cleaners, such as Drano™, and oven cleaners, such as Easy-Off™, which make use of its ability to convert fats to water-soluble soap.

In the process of doing so, however, relatively large amounts of lye may generate enough heat to boil the water in a drain, causing the water to shoot upward. For this reason, it is not advisable to stand near a drain being treated with lye. In a closed oven, this is not a danger, of course; and after the cleaning process is complete, the converted fats (now in the form of soap) can be dissolved and wiped off with a sponge.

WHERE TO LEARN MORE

"Acids and Bases Frequently Asked Questions." General Chemistry Online (Web site). <http://antoine.fsu.umd.edu/chem/senese/101/acidbase/faq.shtml> (June 7, 2001).

"Acids, Bases, and Salts." Chemistry Coach (Web site). <http://www.chemistrycoach.com/acids.htm> (June7, 2001).

"Acids, Bases, and Salts." University of Akron, Department of Chemistry (Web site). <http://ull.chemistry.uakron.edu/genobc/Chapter_09/title.html> (June 7, 2001).

ChemLab. Danbury, CT: Grolier Educational, 1998.

Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

Haines, Gail Kay. What Makes a Lemon Sour? Illustratedby Janet McCaffery. New York: Morrow, 1977.

Oxlade, Chris. Acids and Bases. Chicago: Heinemann Library, 2001.

Patten, J.M. Acids and Bases. Vero Beach, FL: Rourke Book Company, 1995.

Walters, Derek. Chemistry. Illustrated by Denis Bishopand Jim Robins. New York: F. Watts, 1982.

Zumdahl, Steven S. Introductory Chemistry A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.

KEY TERMS

ACID:

A substance that, in its edible form, is sour to the taste, and in non-edible forms, is often capable of dissolving metals. Acids and bases react to form salts and water. These are all phenomenological definitions, however, in contrast to the three structural definitions of acids and bases—the Arrhenius, Brønsted-Lowry, and Lewis acid-base theories.

ALKALI:

A term referring to the soluble hydroxides of the alkali and alkaline earth metals. Once "alkali" was used for the class of substances that react with acids to form salts; today, however, the more general term base is preferred.

ALKALINITY:

An adjectival term used to identify the degree to which a substance displays the properties of a base.

ANION:

The negatively charged ion that results when an atom gains one or more electrons. "Anion" is pronounced "AN-ie-un".

AQUEOUS SOLUTION:

A substance in which water constitutes the solvent. A large number of chemical reactions take place in an aqueous solution.

ARRHENIUS ACID-BASE THEORY:

The first of three structural definitions of acids and bases. Formulated by Swedish chemist Svante Arrhenius (1859-1927), the Arrhenius theory defines acids and bases according to the ions they produce in an aqueous solution: an acid produces hydrogen ions (H⁺), and a base hydroxide ions (OH⁻).

BASE:

A substance that, in its edible form, is bitter to the taste. Bases tend to be slippery to the touch, and in reaction with acids they produce salts and water. Bases and acids are most properly defined, however, not in these phenomenological terms, but by the three structural definitions of acids and bases—the Arrhenius, Brønsted-Lowry, and Lewis acid-base theories.

BRØNSTED-LOWRY ACID-BASE THEORY:

The second of three structural definitions of acids and bases. Formulated by English chemist Thomas Lowry (1874-1936) and Danish chemist J. N. Brønsted (1879-1947), Brønsted-Lowry theory defines an acid as a proton (H⁺) donor, and a base as a proton acceptor.

CATION:

The positively charged ion that results when an atom loses one or more electrons. "Cation" is pronounced "KAT-ie-un".

CHEMICAL SPECIES:

A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.

CONJUGATE ACID:

An acid formed when a base accepts a proton (H⁺).

CONJUGATE ACID-BASE PAIR:

The acid and base produced when an acid donates a single proton to a base. In the reaction that produces this pair, the acid and base switch identities. By donating aproton, the acid becomes a conjugate base, and by receiving the proton, the base becomes a conjugate acid.

CONJUGATE BASE:

A base formed when an acid releases a proton.

ION:

An atom or atoms that has lost or gained one or more electrons, and thus has a net electric charge. There are two types of ions: anions and cations.

IONIC BONDING:

A form of chemical bonding that results from attractions between ions with opposite electric charges.

IONIC COMPOUND:

A compound in which ions are present. Ionic compounds contain at least one metal and non metal joined by an ionic bond.

LEWIS ACID-BASE THEORY:

The third of three structural definitions of acids and bases. Formulated by American chemist Gilbert N. Lewis (1875-1946), Lewis theory defines an acid as the reactant that accepts an electron pair from another reactant in a chemical reaction, and a base as the reactant that donates an electron pair to another reactant.

PH SCALE:

A logarithmic scale for determining the acidity or alkalinity of a substance, from 0 (virtually pure acid) to 7(neutral) to 14 (virtually pure base).

PHENOMENOLOGICAL:

A term describing scientific definitions based purely on experimental phenomena. These convey only part of the picture, however—primarily, the part a chemist can perceive either through measurement or through the senses, such as sight. A structural definition is therefore usually preferable to a phenomenological one.

REACTANT:

A substance that interacts with another substance in a chemical reaction, resulting in the creation of a product.

SALTS:

Ionic compounds formed by the reaction between an acid and a base. In this reaction, one or more of the hydrogenions of an acid is replaced with another positive ion. In addition to producing salts, acid-base reactions produce water.

SOLUTION:

A homogeneous mixture in which one or more substances (thesolute) is dissolved in one or more other substances (the solvent)—for example, sugar dissolved in water.

SOLVENT:

A substance that dissolvesanother, called a solute, in a solution.

STRUCTURAL:

A term describing scientific definitions based on aspects of molecular structure and behavior rather than purely phenomenological data.

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

Organic acids

Uses of acids and bases

Resources

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

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

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.

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

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.

Acids and bases

Acids and bases are chemical compounds that have distinctive properties in water solution. The sour taste of a lemon, lime, or grapefruit, for example, is caused by citric acid. The slippery feel of ammonia, a common base, is characteristic of all bases. One of the most interesting properties of acids and bases is the way they react with indicators. An indicator is a material that changes color in the presence of an acid or a base. For instance, the hydrangea flower can be either pink or blue, depending on the amount of acid or base present in the soil in which it is planted.

Acids and bases have been known since prehistoric times. Vinegar, for example, is a water solution of acetic acid that has been used for centuries. The first modern definitions for acids and bases were suggested by Swedish chemist Svante Arrhenius (1859–1927). Arrhenius proposed that acids be defined as chemicals that produce positively charged hydrogen ions, H⁺, in water. By comparison, he suggested that bases are compounds that produce negatively charged hydroxide ions, OH⁻, in water. Acids and bases react with each other in a reaction called neutralization. In a neutralization reaction, the hydrogen ion from an acid and the hydroxide ion from a base react to form a molecule of water:

H⁺ + OH⁻ → H₂O

Other definitions of acids and bases

Since the time of Arrhenius, chemists have adopted other ways of defining acids and bases. In 1923, English chemist Thomas Lowry (1874–1936) and Danish chemists J. N. Brønsted (1879–1947) and N. Bjerrum (1879–1958) suggested defining acids as chemicals that donate a proton (specifically H⁺) in a chemical reaction and bases as chemicals that accept a proton. This definition is slightly more comprehensive than Arrhenius's definition and, in many cases, more useful to chemists.

Another definition of acids and bases was suggested in 1923 by American chemist Gilbert Newton Lewis (1875–1946). According to Lewis, an acid could be thought of as any compound that accepts a pair of electrons from another substance; bases, on the other hand, could be thought of as compounds that donate a pair of electrons.

Most acids and bases fit all three of these definitions, but some are covered by only one or both of the more modern definitions.

Strong and weak acids and bases

One of the most important characteristics of acids and bases is their strength. The strength of an acid or base depends on the number of hydrogen ions or hydroxide ions produced in water solution. For example, suppose that 100 molecules of an acid are added to water. Of those 100 molecules, imagine that 99 release hydrogen ions. That acid is said to be a strong acid. In comparison, suppose that 100 molecules of a second acid release only 10 hydrogen ions in water. That acid is said to be a weak acid.

Examples of strong acids are hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). Among the best-known strong bases are sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide (CaOH₂). The weak acids include acetic acid (HC₂H₃O₂), lactic acid (CH₃CHOHCOOH), and oxalic acid (H₂C₂O₄). The most familiar weak base is ammonia (NH₃).

Neutralization

Acids and bases are sometimes described as the chemical opposites of each other. If equivalent quantities of an acid and a base are combined, the two compounds react to form a salt and water. For example:

HCl + NaOH → NaCl + H₂O

hydrochloric acid + sodium hydroxide → sodium chloride + water

This process is known as neutralization.

Neutralization has many practical applications. For example, agricultural land is often too acidic or too basic to grow certain crops. Farmers can add either a weak acid or a weak base to produce the level of acidity or basicity in which various plants grow best. Reclamation (restoration) of land once used for mining also involves neutralization reactions. Such land is often too acidic for plants to grow. Treating the land with calcium oxide neutralizes acids remaining in the soil and restores its fertility.

Neutralization is also used to deal with environmental problems. Gases produced in factories, power generating plants, and other industrial facilities are usually acidic. When they escape into the air, they react with water to form acid rain. Scrubbers that contain bases can be attached to the inside of smokestacks in such plants to neutralize acids in escaping gases. Treatment with acids or bases is also used to neutralize hazardous chemicals produced by a variety of manufacturing operations.

Uses of acids and bases

Acids and bases can be found everywhere in the world around us. Lactic acid occurs in sour milk, citric acid in citrus fruits, oxalic acid in rhubarb, malic acid in apples, and tartaric acid in wine. Baking soda, antacids, and lye all contain bases.

Acids and bases are also used widely in industry. Three of the top ten chemicals produced in the United States each year are acids or bases. In 1994, 40 billion kilograms (or about 90 billion pounds) of sulfuric acid were manufactured in the United States, making it the number one chemical in the chemical industry. In addition, 12 billion kilograms (about 26 billion pounds) of sodium hydroxide and 11 billion kilograms (about 25 billion pounds) of phosphoric acid were produced.

The most important single use of acids and bases is in the manufacture of other chemicals. Fertilizers, synthetic fabrics, pigments, petroleum, iron and steel, explosives, dyes, plastics, pesticides, soaps and detergents, paper, film, and many other chemicals are produced from acids and bases. They are also used for various other purposes, including cleaning surfaces, refining oil and sugar, electroplating metals, and treating food products.

Acid-Base Balance

Definition

Acid-base balance can be defined as homeostasis of the body fluids at a normal arterial blood pH ranging between 7.37 and 7.43.

Description

An acid is a substance that acts as a proton donor. In contrast, a base, also known as an alkali, is frequently defined as a substance that combines with a proton to form a chemical bond. Acid solutions have a sour taste and produce a burning sensation with skin contact. A base is any chemical compound that produces hydroxide ions when dissolved in water. Base solutions have a bitter taste and a slippery feel. Despite variations in metabolism, diet, and environmental factors, the body's acid-base balance, fluid volume, and electrolyte concentration are maintained within a narrow range.

Function

Many naturally occurring acids are necessary for life. For example, hydrochloric acid is secreted by the stomach to assist with digestion. The chemical composition of food in the diet can have an effect on the body's acid-base production. Components that affect acid-base balance include protein, chloride, phosphorus, sodium, potassium, calcium, and magnesium. In addition, the rate at which nutrients are absorbed in the intestine will alter acid-base balance.

Cells and body fluids contain acid-base buffers, which help prevent rapid changes in body fluid pH over short periods of time, until the kidneys pulmonary systems can make appropriate adjustments. The kidneys and pulmonary system then work to maintain acid-base balance through excretion in the urine or respiration. The partial pressure of carbon dioxide gas (PCO₂) in the pulmonary system can be measured with a blood sample and correlates with blood carbon dioxide (CO₂) levels. PCO₂ can then be used as an indicator of the concentration of acid in the body. The concentration of base in the body can be determined by measuring plasma bicarbonate (HCO₃⁻) concentration. When the acid-base balance is disturbed, the respiratory system can alter PCO₂ quickly, thus changing the blood pH and correcting imbalances. Excess acid or base is then excreted in the urine by the renal system to control plasma bicarbonate concentration. Changes in respiration occur primarily in minutes to hours, while renal function works to alter blood pH within several days.

Role in human health

Production of CO₂ is a result of normal body metabolism. Exercise or serious infections will increase the production of CO₂ through increased respiration in the lungs. When oxygen (O₂) is inhaled and CO₂ is exhaled, the blood transports these gases to the lungs and body tissues. The body's metabolism produces acids that are buffered and then excreted by the lungs and kidneys to maintain body fluids at a neutral pH. Disruptions in CO₂ levels and HCO₃⁻ create acid-base imbalances. When acid-base imbalances occur, the disturbances can be broadly divided into either acidosis (excess acid) or alkalosis (excess base/alkali).

Common diseases and disorders

Acid-base metabolism imbalances are often characterized in terms of the HCO₃⁻/CO₂ buffer system. Acid-base imbalances result primarily from metabolic or respiratory failures. An increase in HCO₃⁻ is called metabolic alkalosis, while a decrease in the same substance is called metabolic acidosis. An increase in PCO₂, on the other hand, is known as respiratory acidosis, and a decrease in the same substance is called respiratory alkalosis.

Acidosis

Acidosis is a condition resulting from higher than normal acid levels in the body fluids. It is not a disease, but may be an indicator of disease. Metabolic acidosis is related to processes that transform food into energy and body tissues. Conditions such as diabetes, kidney failure, severe diarrhea, and poisoning can result in metabolic acidosis. Mild acidosis is often compensated by the body in a number of ways. However, prolonged acidosis can result in heavy or rapid breathing, weakness, and headache. Acidemia (arterial pH < 7.35) is an accumulation of acids in the bloodstream that may occur with severe acidosis when the acid load exceeds respiratory capacity. This condition can sometimes result in coma and, if the pH falls below 6.80, it will lead to death. Diabetic ketoacidosis is a condition where excessive glucagon and a lack of insulin contribute to the production of ketoacids in the liver. This condition can be caused by chronic alcoholism and poor carbohydrate utilization.

Respiratory acidosis is caused by the lungs's failure to remove excess carbon dioxide from the body, reducing the pH in the body. Several conditions, including chest injury, blockage of the upper air passages, and severe lung disease, may lead to respiratory acidosis. Blockage of the air passages may be caused by bronchitis, asthma, or airway obstruction, resulting in mild or severe acidosis. Regular, consistent retention of carbon dioxide in the lungs is referred to as chronic respiratory acidosis. This disorder results in only mild acidosis because it is balanced by increased bicarbonate production.

The predominant symptoms of acidosis are sometimes difficult to distinguish from symptoms of an underlying disease or disorder. Mild conditions of acidosis may be asymptomatic or may be accompanied by weakness or listlessness, nausea, and vomiting. Most often, severe metabolic acidosis (pH < 7.20) is associated with increased respiration to compensate for a shortage of HCO₃⁻. This is followed by a secondary decrease in PCO₂ that occurs as part of respiratory compensation process. Treatment options for acidosis typically require correction of the underlying condition by venous administration of sodium bicarbonate or another alkaline substance.

Alkalosis

Alkalosis is a condition resulting from a higher than normal level of base/alkali in the body fluids. An excessive loss of HCO₃⁻ in the blood causes metabolic alkalosis. The body can compensate for mild alkalinity, but prolonged alkalosis can result in convulsions, muscular weakness, and even death if the pH rises above 7.80. Alkalosis can be caused by drugs or disorders that upset the normal acid-base balance. Prolonged vomiting and hyperventilation (abnormally fast, deep breathing) can result in alkalosis.

The predominant symptoms of alkalosis are neuromuscular hyperexcitability and irritability. Alkalemia (abnormal blood alkalinity) increases protein binding of ionized calcium even though plasma total calcium does not change. Severe cases may induce hypocalcemia (a low level of plasma calcium). Low plasma potassium leads to a condition called hypokalemic alkalosis. It is frequently accompanied by metabolic alkalosis, resulting in cramping, muscle weakness, polyuria, and ileus (obstruction of the intestines). Diuretic medications may cause hypokalemic alkalosis. Prolonged vomiting may induce hypochloremic alkalosis (a large loss of chloride). The kidneys may conserve bicarbonate in order to compensate for the chloride reduction. Compensated alkalosis results when the body has partially compensated for alkalosis, and has restored normal acid-base balances. However, in compensated alkalosis, abnormal bicarbonate and carbon dioxide levels persist.

KEY TERMS

Acid— (a) Any ionic or molecular substance that can act as a proton donor; (b) A sour-tasting substance, like vinegar; (c) A chemical compound that can react with a base to form a salt.

Acidosis— A dangerous condition where the blood and body tissues are less alkaline (or more acidic) than normal.

Alkalosis— Excessive alkalinity of the blood and body tissue.

Alkalemia— Abnormal blood alkalinity.

Base— (a) Any ionic or molecular substance that can act as a proton acceptor; (b) A bitter-tasting substance which has a soapy feel; (c) A chemical compound that can react with an acid to form a salt. A base can also be called an alkali.

Bicarbonate— A salt of carbonic acid produced by neutralizing a hydrogen ion.

Diabetic ketoacidosis— A condition characterized by excessive thirst and urination. Other symptoms may include appetite loss, nausea, vomiting, and rapid deep breathing.

Diuretic— An agent or drug that eliminates excessive water in the body by increasing the flow of urine.

Electrolyte— A substance such as an acid, bases, or salt. An electrolyte's water solution will conduct an electric current and ionizes. Acids, bases, and salts are electrolytes.

Homeostasis— An organism's regulation of body processes to maintain internal equilibrium in temperature and fluid content.

Hypochloremic alkalosis— A large loss of chloride.

Hypokalemic alkalosis— Low plasma potassium.

pH— The negative logarithm of H+ (hydrogen) concentration.

Alkalosis requires correction of the underlying condition and may involve venous administration of a weak acid to restore normal balance. If the source of alkalosis is excessive drug intake, it may be appropriate to reduce intake to restore the normal acid-base balance.

Respiratory alkalosis results from decreased CO₂ levels caused by conditions such as hyperventilation (a faster breathing rate), anxiety, and fever. The pH is elevated in the body. Hyperventilation causes the body to lose excess carbon dioxide in expired air and can be triggered by altitude or a disease that reduces the amount of oxygen in the blood. Symptoms of respiratory alkalosis may include dizziness, lightheadedness, and numbing of the hands and feet. Treatments include breathing into a paper bag or a mask that induces rebreathing of carbon dioxide.

Resources

BOOKS

Shaw, Patricia, ed. Fluids & Electrolytes Made Incredibly Easy! Springhouse, PA: Springhouse Publishing Co., 1997.

PERIODICALS

Remer, T. "Influence of diet on acid-base balance." Seminars in Dialysis (2000) 13, no. 4: 221-226.

OTHER

Bookallil, Michael. "pH of the blood: acid-base balance." 〈http://www.usyd.edu.au/su/anaes/lectures/acidbase_mjb/frameversion.html〉.

Grogono, Alan. "Acid-base tutorial." 〈http://www.acidbase.com〉.