To an extent, acids and bases can be defined in terms of factors that are apparent to the senses: edible acids taste sour, for instance, while bases are bitter-tasting and slippery to the touch. The best way to understand these two types of substances, however, is in terms of their behavior in chemical reactions. Not only do the reactions of acids and bases result in the creation of salts and water, but acids and bases can be defined by the ways in which they participate in a reaction—for instance, by donating or accepting electron pairs. The reaction of acids and bases to form water and salts is called neutralization, and it has a wide range of applications, including the promotion of plant growth in soil and the treatment of heartburn in the human stomach. Neutralization also makes it possible to test substances for their pH level, a measure of the degree to which the substance is acidic or alkaline.
HOW IT WORKS
Phenomenological Definitions of Acids and Bases
Before studying the reactions of acids and bases, it is necessary to define exactly what each is. This is not as easy as it sounds, and the Acids and Bases essay discusses in detail a subject covered more briefly here: the arduous task chemists faced in developing a workable distinction. Let us start with the phenomenological differences between the two—that is, aspects relating to things that can readily be observed without referring to the molecule properties and behaviors of acids and bases.
Acids are fairly easy to understand on the phenomenological level: the name comes from the Latin term acidus, or "sour," and many sour substances from daily life—lemons, for instance, or vinegar—are indeed highly acidic. In fact, lemons and most citrus fruits contain citric acid (C6H8O7), while the acidic quality of vinegar comes from acetic acid (CH3COOH). In addition, acids produce characteristic colors in certain vegetable dyes, such as those used in making litmus paper.
The word "base," as it is used in this context, may be a bit more difficult to appreciate on a sensory level. It helps, perhaps, if the older term "alkali" is used, though even so, people tend to think of alkaline substances primarily in contrast to acids. "Alkali," which serves to indicate the basic quality of both the alkali metal and alkaline earth metal families of elements, comes from the Arabic al-qili. The latter refers to the ashes of the seawort plant, which usually grows in marshy areas and, in the past, was often burned to produce soda ash for making soap.
The reason chemists of today use the word "base" instead of "alkali" is that the latter term has a narrower meaning: all alkalies are bases, but not all bases are alkalies. Originally referring only to the ashes of burned plants containing either sodium or potassium, alkali was eventually used to designate the soluble hydroxides of the alkali and alkaline earth metals. Among these are sodium hydroxide or lye; magnesium hydroxide (found in milk of magnesia); potassium hydroxide, which appears in soaps; and other compounds. Because these represent only a few of the substances that react with acids in the ways discussed in this essay, the term "base" is preferred.
THE FORMATION OF SALTS.
As chemistry evolved, and physical scientists became aware of the atomic and molecular substructures that make up the material world, they developed more fundamental distinctions between acids and bases. By the early twentieth century, chemists had applied structural distinctions between acids and bases—that is, definitions based on the molecular structures and behaviors of those substances.
An important intermediary step occurred as chemists came to the conclusion that reactions of acids and bases form salts and water. For instance, in an aqueous solution, hydrochloric acid or HCl(aq ) reacts with the base sodium hydroxide, designated as NaOH(aq ), to form sodium chloride, or common table salt (NaCl[aq ]) and H2O. What happens is that the sodium (Na) ion (an atom with an electric charge) in sodium hydroxide switches places with the hydrogen ion in hydrochloric acid, resulting in the creation of NaCl and water.
Ions themselves had yet to be defined in 1803, when the great Swedish chemist Jons Berzelius (1779-1848) added another piece to the foundation for a structural definition. Acids and bases, he suggested, have opposite electric charges. In this, he was about eight decades ahead of his time: only in 1884 did his countryman Svante Arrhenius (1859-1927) introduce the concept of the ion. This, in turn, enabled Arrhenius to formulate the first structural distinction between acids and bases.
The Arrhenius Acid-Base Theory
Arrhenius acid-base theory defines the two substances with regard to their behavior in an aqueous solution: an acid is any compound that produces hydrogen ions (H+), and a base is one that produces hydroxide ions (OH−) when dissolved in water. This occurred, for instance, in the reaction discussed above: the hydrochloric acid produced a hydrogen ion, while the sodium hydroxide produced a hydroxide ion, and these two ions bonded to form water.
Though it was a good start, Arrhenius's theory was limited to reactions in aqueous solutions. In addition, it confined its definition of acids and bases only to those ionic compounds, such as hydrochloric acid or sodium hydroxide, that produced either hydrogen or hydroxide ions. But ammonia, or NH3, acts like a base in aqueous solutions, even though it does not produce the hydroxide ion. These shortcomings pointed to the need for a more comprehensive theory, which came with the formulation of the Brønsted-Lowry definition.
The BrØnsted-Lowry Acid-Base Theory
Developed by English chemist Thomas Lowry (1874-1936) and Danish chemist J. N. Brønsted (1879-1947), 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+, a cation (positively charged ion) of hydrogen.
Elemental hydrogen, called protium to distinguish it from its isotopes, has just one proton and one electron—no neutrons. Therefore, the hydrogen cation, which has to lose its sole electron to gain a positive charge, is essentially nothing but a proton. It is thus at once an atom, an ion, and a proton, but the ionization of hydrogen constitutes the only case in which this is possible.
Thus when the term "proton donor" or "proton acceptor" is used, it does not mean that a proton is splitting off from an atom or joining another, as in a nuclear reaction. Rather, when an acid behaves as a proton donor, this means that the hydrogen proton/ion/atom is separating from an acidic compound; conversely, when a base acts as a proton acceptor, the positively charged hydrogen ion is bonding with the basic compound.
REACTIONS IN BRØNSTED-LOWRY ACID-BASE THEORY.
In representing Brønsted-Lowry acids and bases, the symbols HA and A−, respectively, are used. These appear in the equation representing the most fundamental type of Brønsted-Lowry acid-base reaction: HA(aq ) + H2O(l ) →H3O+(aq ) + A−(aq ). The symbols (aq ), (l ), and →are explained in the Chemical Reactions essay. In plain English, this equation states that when an acid in an aqueous solution reacts with liquid water, the result is the creation of H3O+, known as the hydronium ion, along with a base. Both products of the reaction are dissolved in an aqueous solution.
Because water molecules are polar, 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 the hydronium ion.
The hydronium ion, in this equation, 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 includes all Arrhenius acids and bases, as well as other chemical species not encompassed in Arrhenius theory. As mentioned earlier, ammonia is a base, yet it does not produce OH− ions; however, it does accept a proton from a water molecule. Water can serve either as an acid or base; in this instance, it is an acid, and in reaction with ammonia, it produces the conjugate acid-base pair of NH4+ (an ammonium ion) and OH−. Ammonia did not produce the hydroxide ion here; rather, OH− is the conjugate base that resulted when the water molecule lost its H+ atom (i.e., a proton.)
The Lewis Acid-Base Theory
The Brønsted-Lowry model still had its limitations, in that it only described compounds containing hydrogen. American chemist Gilbert N. Lewis (1875-1946), however, developed a theory of acids and bases that makes no reference to the presence of hydrogen. Instead, it relates to something much more fundamental: the fact that chemical bonding always involves pairs of electrons.
Lewis acid-base theory defines an acid as 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. Note that, as with the Brønsted-Lowry definition, the Lewis definition is reaction-dependant. Instead of defining a compound as an acid or base in its own right, it identifies these in terms of how the compound reacts with another.
The Lewis definition encompasses all the situations covered by the others, as well as many other reactions not described in the theories of either Arrhenius or Brønsted-Lowry. In particular, Lewis theory can be used to differentiate the acid and base in chemical reactions where ions are not produced, something that takes it far beyond the scope of Arrhenius theory. Also, Lewis theory addresses situations in which there is no proton donor or acceptor, thus offering an improvement over Brønsted-Lowry.
When boron trifluoride (BF3) and ammonia (NH3), both in the gas phases, react to produce boron trifluoride ammonia complex (F3BNH3), boron trifluoride accepts an electron pair. Therefore, it is a Lewis acid, while ammonia—which donates the electron pair—can be defined as a Lewis base. This particular reaction involves hydrogen, but since the operative factor in Lewis theory relates to electron pairs and not hydrogen, the theory can be used to address reactions in which that element is not present.
Dissociation is the separation of a molecule into ions, and it is a key factor for evaluating the "strength" of acids and bases. The more a substance is prone to dissociation, the better it can conduct an electric current, because the separation of charges provides a "pathway" for the current's flow. A substance that dissociates completely, or almost completely, is called a strong electrolyte, whereas one that dissociates only slightly (or not at all) is designated as a weak electrolyte.
The terms "weak" and "strong" are also applied to acids and bases. For instance, vinegar is a weak acid, because it dissociates only slightly, and therefore conducts little electric current. By contrast, hydrochloric acid (HCl) is a strong acid, because it dissociates almost completely into positively charged hydrogen ions and negatively charged chlorine ones. Represented symbolically, this is: HCl →H+ + Cl−.
A REACTION INVOLVING A STRONG ACID.
It may seem a bit backward that a strong acid or base is one that "falls apart," while the weak one stays together. To understand the difference better, let us return to the reaction described earlier, in which an acid in aqueous solution reacts with water to produce a base in aqueous solution, along with hydronium: HA(aq ) + H2O(l ) →H3O+(aq ) + A−(aq ). Instead of using the generic symbols HA and A−, however, let us substitute hydrochloric acid (HCl) and chloride (Cl−) respectively.
The reaction HCl(aq ) + H2O(l ) →H3O+(aq ) + Cl−(aq ) is a reversible one, and for that reason, the symbol for chemical equilibrium (⇋) can be inserted in place of the arrow pointing to the right. In other words, the substances on the right can just as easily react, producing the substances on the left. In this reverse reaction, the reactants of the forward reaction would become products, and the products of the forward reaction serve as the reactants.
However, the reaction described here is not perfectly reversible, and in fact the most proper chemical symbolism would show a longer arrow pointing to the right, with a shorter arrow pointing to the left. Due to the presence of a strong electrolyte, there is more forward "thrust" to this reaction.
Because it is a strong acid, the hydrogen chloride in solution is not a set of molecules, but a collection of H+ and Cl− ions. In the reaction, the weak Cl− ions to the right side of the equilibrium symbol exert very little attraction for the H+ ions. Instead of bonding with the chloride, these hydrogen ions join the water (a stronger base) to form hydronium.
The chloride, incidentally, is the conjugate base of the hydrochloric acid, and this illustrates another principal regarding the "strength" of electrolytes: a strong acid produces a relatively weak conjugate base. Likewise, a strong base produces a relatively weak conjugate acid.
THE STRONG ACIDS AND BASES.
There are only a few strong acids and bases, which are listed below:
- Hydrobromic acid (HBr)
- Hydrochloric acid (HCl)
- Hydroiodic acid (HI)
- Nitric acid (HNO3)
- Perchloric acid (HClO4)
- Sulfuric acid (H2SO4)
- Barium hydroxide (Ba[OH]2)
- Calcium hydroxide (Ca[OH]2)
- Lithium hydroxide (LiOH)
- Potassium hydroxide (KOH)
- Sodium hydroxide (NaOH)
- Strontium hydroxide (Sr[OH]2)
Virtually all others are weak acids or bases, meaning that only a small percentage of molecules in these substances ionize by dissociation. The concentrations of the chemical species involved in the dissociation of weak acids and bases are mathematically governed by the equilibrium constant K..
Neutralization is the process whereby an acid and base react with one another to form a salt and water. The simplest example of this occurs in the reaction discussed earlier, in which hydrochloric acid or HCl(aq ) reacts with the base sodium hydroxide, designated as NaOH(aq ), in an aqueous solution. The result is sodium chloride, or common table salt (NaCl[aq ]) and H2O. This equation is written thus: HCl(aq ) + NaOH(aq ) →NaCl(aq ) + H2O.
The human stomach produces hydrochloric acid, commonly known as "stomach acid." It is generated in the digestion process, but when a person eats something requiring the stomach to work overtime in digesting it—say, a pizza—the stomach may generate excess hydrochloric acid, and the result is "heartburn." When this happens, people often take antacids, which contain a base such as aluminum hydroxide (Al[OH]3) or magnesium hydroxide (Mg[OH]2).
When a person takes an antacid, the reaction leads to the creation of a salt, but not the salt with which most people are familiar—NaCl. As shown above, that particular salt is the product of a reaction between hydrochloric acid and sodium hydroxide, but a person who ingested sodium hydroxide (a substance used to unclog drains and clean ovens) would have much worse heartburn than before! In any case, the antacid reacts with the stomach acid to produce a salt, as well as water, and thus the acid is neutralized.
When land formerly used for mining is reclaimed, the acidic water in the area must be neutralized, and the use of calcium oxide (CaO) as a base is one means of doing so. Acidic soil, too, can be neutralized by the introduction of calcium carbonate (CaCO3) or limestone, along with magnesium carbonate (MgCO3). If soil is too basic, as for instance in areas where there has been too little precipitation, acid-like substances such as calcium sulfate or gypsum (SaSO4) can be used. In either case, neutralization promotes plant growth.
TITRATION AND pH.
One of the most important applications of neutralization is in titration, the use of a chemical reaction to determine the amount of a chemical substance in a sample of unknown purity. In a typical form of neutralization titration, a measured amount of an acid is added to a solution containing an unknown amount of a base. Once enough of the acid has been added to neutralize the base, it is possible to determine how much base exists in the solution.
Titration can also be used to measure pH ("power of hydrogen") level by using an acid-base indicator. The pH scale assigns values ranging from 0 (a virtually pure acid) to 14 (a virtually pure base), with 7 indicating a neutral substance. An acid-base indicator such as litmus paper changes color when it neutralizes the solution.
The transition interval (the pH at which the color of an indicator changes) is different for different types of indicators, and thus various indicators are used to measure substances in specific pH ranges. For instance, methyl red changes from red to yellow across a pH range of 4.4 to 6.2, so it is most useful for testing a substance suspected of being moderately acidic.
A buffered solution is one that resists a change in pH even when a strong acid or base is added to it. This buffering results from the presence of a weak acid and a strong conjugate base, and it can be very important to organisms whose cells can endure changes only within a limited range of pH values. Human blood, for instance, contains buffering systems, because it needs to be at pH levels between 7.35 and 7.45.
The carbonic acid-bicarbonate buffer system is used to control the pH of blood. The most important chemical equilibria (that is, reactions involving chemical equilibrium) for this system are: H+ + HCO3− ⇋ H2CO3 ⇋ H20 + CO2. In other words, the hydrogen ion (H+) reacts with the hydrogen carbonate ion (HCO3−) to produce carbonic acid (H2CO3). The latter is in equilibrium with the first set of reactants, as well as with water and carbon dioxide in the forward reaction.
The controls the pH level by changing the concentration of carbon dioxide by exhalation. In accordance with Le Châtelier's principle, this shifts the equilibrium to the right, consuming H+ ions. In normal blood plasma, the concentration of HCO3− is about 20 times as great as that of H2CO3, and this large concentration of hydrogen carbonate gives the buffer a high capacity to neutralize additional acid. The buffer has a much lower capacity to neutralize bases because of the much smaller concentration of carbonic acid.
Water: Both Acid and Base
Water is an amphoteric substance; in other words, it can serve either as an acid or a base. When water experiences ionization, one water molecule serves as a Brønsted-Lowry acid, donating a proton to another water molecule—the Brønsted-Lowry base. This results in the production of a hydroxide ion and a hydronium ion: H2O(l ) + H2O(l ) ⇋ H3O+(aq ) + OH−(aq ).
This equilibrium equation is actually one in which the tendency toward the reverse reaction is much greater; therefore the equilibrium symbol, if rendered in its most proper form, would show a much shorter arrow pointing toward the right. In water purified by distillation, the concentrations of hydronium (H3O+) and hydroxide (OH−) ions are equal. When multiplied by one another, these yield the constant figure 1.0 · 10−14, which is the equilibrium constant for water. In fact, this constant—denoted as K w—is called the ion-product constant for water.
Because the product of these two concentrations is always the same, this means that if one of them goes up, the other one must go down in order to yield the same constant. This explains the fact, noted earlier, that water can serve either as an acid or base—or, if the concentrations of hydronium and hydroxide ions are equal—as a neutral substance. In situations where the concentration of hydronium is higher, and the hydroxide concentration automatically decreases, water serves as an acid. Conversely, when the hydroxide concentration is high, the hydronium concentration decreases correspondingly, and the water is a base.
WHERE TO LEARN MORE
"Acids and Bases." Vision Learning (Web site). <http://www.visionlearning.com/library/science/chemistry-2/CHE2.2-acid_base.htm> (June 7, 2001).
"Acids, Bases, and Chemical Reactions" Open Access College (Web site). <http://oac.schools.sa.edu.au/8-10science/acids.htm> (June 7, 2001).
"Acids, Bases, pH." About.com (Web site). <http://chemistry.about.com/science/chemistry/cs/acidsbasesph/ (June 7, 2001).
"Acids, Bases, and Salts" (Web site). <http://edie.cprost.sfu.ca/~rhlogan/ionic_eq.html> (June 7, 2001).
"Junior Part: Acids, Bases, and Salts" (Web site). <http://www.rjclarkson.demon.co.uk/junior/junior4.htm> (June 7, 2001).
Knapp, Brian J. Acids, Bases, and Salts. Danbury, CT: Grolier Educational, 1998.
Moje, Steven W. Cool Chemistry: Great Experiments with Simple Stuff. New York: Sterling Publishing Company, 1999.
Walters, Derek. Chemistry. Illustrated by Denis Bishop and Jim Robins. New York: F. Watts, 1982.
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.
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.
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.
An adjectival term used to identify the degree to which a substance displays the properties of a base.
A term describing a substance that can serve either as an acid or a base. Water is the most significant amphoteric substance.
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 anaqueous solution an acid produces hydrogen ions (H+), and a base hydroxide ions (OH−).
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.
In the context of acids and bases, the word is the counterpart to "acidic," identifying the base-like quality of a substance.
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.
A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.
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.
A base formed when an acid releases a proton.
The separation of molecules into ions.
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.
A form of chemical bonding that results from attractions between ions with opposite electric charges.
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.
The process whereby an acid and base react with one another to form a salt and water.
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).
A term describing scientific definitions based purely on experimental phenomena. These only convey 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.
A substance that interacts with another substance in a chemical reaction, resulting in the creation of a product.
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.
A homogeneous mixture in which one or more substances (thesolute) is dissolved in another substance (the solvent)—for example, sugar dissolved in water.
A substance that dissolvesanother, called a solute, in a solution.
A substance highly prone to dissociation. The terms "strong acid" or "strong base" refers to those acids or bases which readily dissociate.
A term describing scientific definitions based on aspects of molecular structure and behavior rather than purely phenomenological data.
The use of a chemical reaction to determine the amount of a chemical substance in a sample of unknown purity. Testing pH levels is an example of titration.
The pH level at which the color of an acid-base indicator changes.
A substancethat experiences little or no dissociation. The terms "weak acid" or "weak base" refer to those acids or bases not prone to dissociation.
"Acid-Base Reactions." Science of Everyday Things. . Encyclopedia.com. (January 16, 2019). https://www.encyclopedia.com/science/news-wires-white-papers-and-books/acid-base-reactions
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Acids and bases have been known by their properties since the early days of experimental chemistry. The word "acid" comes from the Latin acidus, meaning "sour" or "tart," since water solutions of acids have a sour or tart taste. Lemons, grapefruit, and limes taste sour because they contain citric acid and ascorbic acid (vitamin C). Another common acid is vinegar, which is the sour liquid produced when apple cider, grape juice, or other plant juices ferment beyond the formation of alcohol. Vinegar is a 5 percent water solution of acetic acid. Besides having a sour taste, acids react with active metals to give hydrogen, they change the colors of indicators (for example, litmus turns from blue to red), and they neutralize bases. Bases change the colors of indicators (litmus turns from red to blue) and they neutralize acids. Hence, bases are considered the chemical opposite of acids.
Most common acid-base reactions take place in water solutions (commonly referred to as aqueous solutions ). One of the earliest definitions of acids, advanced by the Swedish physicist and chemist Svante Arrhenius in 1887, stated that acid ionizes in aqueous solution to produce hydrogen ions (which are protons), H+, and anions ; and a base ionizes in aqueous solution to produce hydroxide ions (OH−) and cations. Later studies of aqueous solutions provided evidence of a small, positively charged hydrogen ion combining with a water molecule to form a hydrated proton, H+(H2O) or H3O+, which is called the hydronium ion. Often, the hydronium ion or hydrated proton is represented as H+ (aq ). Hydrogen chloride (HCl), a gas, is an acid because it dissolves in water to yield hydrogen ions and chloride ions. This water solution of HCl is referred to as hydrochloric acid.
A typical base, according to the Arrhenius definition, is sodium hydroxide (NaOH). It dissolves in water to give sodium ions and hydroxide ions.
In the reaction of an acid with a base in aqueous solution, the hydrogen ions of the acid react with the hydroxide ions of the base to give water. The second product is a salt, which is composed of the positive metal ion from the base and the negative ion from the acid. For example,
HCl (aq ) + KOH (aq ) → H2O (l ) + KCl (aq ) (3)
Since HCl (aq ) and KOH (aq ) are fully ionized in solution, the preceding equation can be written as
H+ (aq ) + Cl− (aq ) K+ (aq ) + OH− (aq ) → H2O (l ) K+ (aq ) + Cl− (aq ) (4)
Ions common to both sides can be canceled to yield
H+ (aq ) + OH− (aq ) → H2O (l ) (5)
This is referred to as the net ionic equation for the neutralization reaction. If H3O+ is substituted for H+ (aq ), the neutralization equation becomes
H3O+ (aq ) + OH− (aq ) → 2 H2O (l ) (6)
Strengths of Acids and Bases
The strength of an acid or base is determined by the extent of its ionization in aqueous solution. Strong acids, such as hydrochloric acid, are 100 percent ionized in aqueous solution, whereas weak acids, such as acetic acid, are less than 5 percent ionized. Experimentally, the extent of ionization is determined by measuring the electrical conductance of solutions. Strong acids and bases are strong electrolytes, and weak acids and bases are weak
|COMMON ACIDS AND BASES|
|Strong Acids||Strong Bases|
|HCl||hydrochloric acid||NaOH||sodium hydroxide|
|HNO3||nitric acid||KOH||potassium hydroxide|
|H2SO4||sulfuric acid||Ba(OH)2||barium hydroxide|
|Weak Acids||Weak Bases|
|H2CO3||carbonic acid||CH2NH2||methyl amine|
electrolytes. Table 1 lists some common acids and bases and indicates whether they are strong or weak.
For weak acids and bases, partial ionization is a dynamic equilibrium between unionized molecules and its ion, as indicated by the double arrow in equation (7). For example, acetic acid is only partially ionized in aqueous solution
CH3COOH (aq ) ⇆ H+ (aq ) + CH3COO− (aq ) (7)
In acetic acid, hydrogen ions and acetate ions recombine to form acetic acid molecules. The double arrow signifies that at any given instant, less than 5 percent of acetic acid molecules dissociate into hydrogen ions and acetate ions, while the hydrogen ions and acetate ions recombine to form acetic acid molecules.
Ammonia (NH3) is a weak base, and although it does not have OH− ions in its formula, it produces the ion on reaction with water.
NH3 (aq ) + H2O (l ) ⇆ NH4+ (aq ) + OH− (aq ) (8)
A major problem with Arrhenius's acid-base theory is that some substances, like ammonia, produce basic solutions and react with acids, but do not contain hydroxide ions. In 1923 Johannes Brønsted, a Danish chemist, and Thomas Lowry, an English chemist, independently proposed a new way to define acids and bases. An acid donates hydrogen ions (also called a proton donor); a base accepts hydrogen ions (also called a proton acceptor). These definitions not only explain all the acids and bases covered by Arrhenius's theory, they also explain the basicity of ammonia and ions such as carbonate, CO32−, and phosphate, PO43.
The Brønsted-Lowry theory includes water as a reactant and considers its acidity or basicity in the reaction. In the partial ionization of acetic acid, water is a base because it accepts the hydrogen ion to form hydronium ion.
THOMAS M. LOWRY (1874–1936)
A meticulous experimenter, Thomas Lowry is best known for his conceptualization of acid–base chemistry. Studies of nitrogenous compounds led Lowry to question fundamental aspects of the role of hydrogen during acid–base reactions. Three months before Brønsted published his theory, Lowry released his own similar thoughts on proton acceptors and donors in print.
In the reaction, a new acid and a new base are formed, which are called the conjugate acid and conjugate base, respectively. The hydronium ion, H3O+, is the conjugate acid of the base, H2O, and the acetate ion, CH3COO−, is
|CONJUGATE ACID-BASE PAIRS|
|Strong acids||H2SO4||HSO4−||Weak bases|
|Weak acids||H2O||OH−||Strong bases|
the conjugate base of acetic acid, CH3COOH. A pair of molecules or ions related to one another by the gain or loss of a single hydrogen ion is called a conjugate acid-base pair. In the reaction of ammonia, water is an acid because it donates a hydrogen ion to ammonia.
This ability of water to donate or accept hydrogen ions, depending on whether it reacts with a base or an acid, is referred to as "amphiprotic." The conjugate acid-base pairs in this reaction are NH3/NH4+ and H2O/OH−.
The Brønsted-Lowry definitions also explain why carbonate salts such as sodium carbonate (washing soda) dissolve in water to give basic solutions. Carbonate ion removes a hydrogen ion from a water molecule, which leaves behind a hydroxide ion:
In the preceding reaction, water and hydroxide ion are a conjugate acid-base pair, whereas carbonate ion and bicarbonate ion are a conjugate base-acid pair. Every Brønsted-Lowry acid has a conjugate base, and every Brønsted-Lowry base has a conjugate acid. Familiarity with conjugate acid-base pairs is important to understanding the relative strengths of acids and bases. Table 2 lists some conjugate acid-base pairs and their relative strengths. Strong acids have weak conjugate bases, and weak acids have strong conjugate bases.
Several common acids have more than one ionizable hydrogen ion (Table1). Each successive hydrogen ion in these polyprotic acids ionizes less readily. For example, sulfuric acid is a strong acid because of the complete ionization of the first hydrogen ion.
H2SO4 (aq ) + H2O (l ) → H3O+ (aq ) + HSO4− (aq ) (12)
The HSO4− also acts as an acid, but it is not 100 percent ionized, so HSO4− is an acid of moderate strength. For example, sodium hydrogen sulfate is used to increase the acidity of swimming pools, whereas sodium carbonate is used to increase the basicity of swimming pools.
HSO4− (aq ) + H2O (l ) ⇆ H3O+ (aq ) + SO42− (aq ) (13)
Phosphoric acid has three ionizable hydrogen ions. Each stepwise ionization of phosphoric acid occurs to a lesser extent than the one before it. Phosphoric acid is stronger than acetic acid because the first step ionizes to a greater extent than acetic acid.
H3PO4 (aq ) + H2O (l ) ⇆ H3O+ (aq ) + H2PO4− (aq ) (14)
However, H2PO4− is a weaker acid than acetic acid because the second ionization is much smaller (by a factor of 105) than the first step.
H2PO4− (aq ) + H2O (l ) ⇆ H3O+ (aq ) + HPO42− (aq ) (15)
The third ionization is also much smaller than the second step (by a factor of 105).
HPO42− (aq ) + H2O (l ) ⇆ H3O+ (aq ) + PO43− (aq ) (16)
The anions of phosphoric acid can also accept hydrogen ions and act as bases with a strong acid like hydrochloric acid. For example,
PO43− (aq ) + H3O+ (aq ) ⇆ HPO42− (aq ) + H2O (l ) (17)
Ions such as H2PO4−, HPO42−, HCO3−, and HSO4− can act as an acid by donating a hydrogen ion or as a base by accepting a hydrogen ion. This ability to act as either an acid or a base is referred to as an amphoteric property.
The pH Scale
The Brønsted-Lowry acid-base definitions are based on the amphiprotic properties of water: Water is capable of acting as both a hydrogen ion donor and a hydrogen ion acceptor, depending on the acidic or basic properties of the dissolved substance (equations 9 and 10). Water can also act as a proton donor and proton acceptor towards itself. This is referred to as the autoionization of water.
H2O (l ) + H2O (l ) ⇆ H3O+ (aq ) + OH− (aq ) (18)
Pure water is neutral because it contains equal numbers of hydronium ions and hydroxide ions. However, pure water only slightly ionizes, about 1 in every 55,000,000 water molecules is ionized at any given time. The actual molar concentration of hydronium ions and hydroxide ions in pure water at 25°C is 1.0 × 10−7. The product of the molarity of the hydronium ions and hydroxide ions of pure water is (1.0 × 10−7) (1.0 × 10−7) = 1.0 × 10−14.
The value of 1.0 × 10−14 is important to the study of aqueous solutions of acids and bases because it is a constant that is always the product of the molar concentration of H3O+ and OH−.
[H3O+] [OH−] = 1.0 × 10−14 (19)
If acid is added to pure water, the concentration of H3O+ will be greater than 1.0 × 10−7, and then the concentration of OH− will be less than 1.0 × 10−7. However, the product of the two must equal 1.0 × 10−14. This relationship is the basis for calculating the concentration of one of the two ions, hydronium or hydroxide, when the other one is known. For example, a 0.1 M solution of hydrochloric acid is 0.1 M in H3O+ since hydrochloric
|pH OF COMMON SUBSTANCES|
|0||1.0||Battery acid, 1 M sulfuric acid|
|1||0.1||Stomach acid, 0.1 M hydrochloric acid|
|2||1 × 10−2||Lemon juice|
|3||1 × 10−3||Vinegar|
|4||1 × 10−4||Soft drink|
|5||1 × 10−5||Rain water|
|6||1 × 10−6||Milk|
|7||1 × 10−7||Pure water|
|8||1 × 10−8||Baking soda, NaHCO3|
|9||1 × 10−9||Washing soda, Na2CO3|
|10||1 × 10−10||Milk of magnesia, Mg(OH)2|
|11||1 × 10−11||Aqueous household ammonia, NH3|
|12||1 × 10−12||Limewater, Ca(OH)2|
|13||1 × 10−13||Drano, 0.1 M NaOH|
|14||1 × 10−14||Drano, 1.0 M NaOH|
acid is fully ionized. From the equation, the molar concentration of OH− is 10−13. For a 0.1 M solution of NaOH, the OH− is 0.1 M, but the hydrogen ion concentration is 10−13. Hence, the value of the exponent for hydronium ion concentration goes from −1 in strong 0.1 M acid to −13 in strong 0.1 M base.
In 1909 the Danish biochemist S. P. L. Sørensen proposed that these exponents be used as a measure of acidity. He devised a scale that would be useful in testing the acidity of Danish beer. Sorensen's scale came to be known as the pH scale, from the French pouvoir hydrogene, which means hydrogen power. pH is defined as the negative logarithm (log) of the hydronium ion concentration.
pH = −log[H3O+] (20)
The brackets around hydronium ion mean moles per liter of hydronium ions.
The pH scale includes values between 0 and 14. The pH of pure water is 7 because [H3O+ ] is 1.0 × 10−7. The pH of 0.1 M HCl is 1[−log 10−1 = −(−1)]. The pH of 0.1 M NaOH is 13(−log 10−13) = 13. The pH scale does not apply to concentrations greater than 1.0 M for a strong acid (pH = 0) or 1.0 M for a strong base (pH = 14).
For solutions in which [H3O+] is not an exact power of 10 (0.1, 0.01,...), a calculator can be used to determine the logarithm. For example, if the [H3O+] is 1.5 × 10−3 M, the logarithm is −3 + log 1.5 = −3 + 0.18 = −2.82, and the pH is −(−2.82) or 2.82. Table 3 provides the pH values of some common solutions.
Many natural substances are acid-base indicators. The most familiar one is litmus, an organic dye extracted from certain lichens. Litmus turns from blue to red in acidic solutions (< pH 7) and from red to blue in basic solutions (∼ pH 7). Some other natural indicators include red cabbage extract, blueberry juice, black tea, beet juice, rhubarb, and tomato leaves, and flowers such as the rose, daylily, blue iris, and purple dahlia. Red cabbage extract undergoes sharp changes of color at several pH values. The deep purple color of red cabbage leaves is caused by a mixture of water-soluble
|Indicator||pH Range||Color Change|
|Thymol blue||1.2 – 2.8||red → yellow|
|Methyl red||4.4 – 6.2||red → yellow|
|Litmus||5 – 8||red → blue|
|Bromothymol blue||6.2 – 7.6||yellow → blue|
|Phenolphthalein||8.0 – 10.0||colorless → pink|
anthocyanins. Over the pH range of 2 to 12, these anthocyanins change from red (pH 2) to pink (pH 4) to purple (pH 6−7) to green (pH 10) to yellow (pH 12), which makes red cabbage extract a "universal indicator."
Acid-base indicators are weak acids and bases. A typical indicator will ionize in aqueous solution according to the equation
HIn (aq ) + H2O (l ) ⇆ H3O+ (aq ) + In− (aq ) (21)
The chemical species HIn and In− are different colors. When the solution is acidic to the degree that the HIn species dominates, it will be the color of HIn. When the solution is more basic with In− dominating, it will be the color of In−. Some common indicators and the pH ranges for their color changes are listed in Table 4.
Buffer solutions contain a base and an acid that can react with an added acid or base, respectively, and they maintain a pH very close to the original value. Buffers usually consist of approximately equal quantities of a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, a buffer solution of acetic acid and its conjugate base, the acetate ion, can neutralize small amounts of a strong acid or strong base as follows:
CH3COOH (aq ) + OH− (aq ) → CH3COO− (aq ) + H2O (l ) (22)
CH3COO− (aq ) + H3O+ (aq ) → CH3COOH (aq ) + H2O (l ) (23)
As illustrated in equations (22) and (23), the addition of either a strong base or a strong acid produces one of the components of the buffer mixture and so the pH does not change. Buffers are limited in their buffer capacity, that is, the amount of a strong acid or strong base that can be added before the pH changes by 1 pH unit.
Buffers are very important to many industrial and natural processes. For example, controlling the pH of blood is essential to human health. The pH of blood is normally 7.40 ± 0.05, and good health depends on the ability of buffers to maintain the pH of blood within this narrow range. If the pH falls below 7.35, a condition known as acidosis occurs; increasing pH above 7.45 leads to alkalosis. Both these conditions can be life threatening. Two buffer systems, H2CO3/HCO3− and H2PO4−/HPO42−, control the pH of the blood.
Lewis Acid-Base Theory
In the early 1930s Gilbert Lewis, an American chemist, proposed a more general acid-base theory that is based on sharing electron pairs rather than proton transfers. A Lewis acid is a substance that can accept a pair of electrons to form a new bond, and a Lewis base is a substance that can donate a pair of electrons to form a new bond. All Arrhenius and Brønsted-Lowry acids and bases are Lewis acids and bases. However, Lewis acid-base theory is more general because a Lewis base can donate an electron pair to something other than H+. For example, the gas phase reaction of NH3 with BF3 is a Lewis acid-base reaction.
Solvent System Acid-Base Theory
Another acid-base theory that is useful for solvents other than water was postulated by American chemist Edward Franklin in 1905. It makes use of the autoionization of solvents, and defines an acid as a solute that produces the positively charged species of the solvent and a base as a solute that produces the negatively charged species of the solvent. In the case of the autoionization of water (equation 18) H3O+ is the acid and OH− is the base. For the nonaqueous solvent, liquid ammonia, the autoionization gives
NH3 (l ) + NH3 (l ) ⇆ NH4+ + NH2− (25)
so an acid in liquid ammonia is any solute that produces NH4+ and a base in liquid ammonia is any solute that produces NH2−. An example of an acid-base reaction in liquid ammonia is
Note that liquid ammonia still falls within the Brønsted-Lowry definitions since NH4+ is a proton donor and NH2− is a proton acceptor.
The Brønsted-Lowry theory, which defines acids as proton donors and bases as proton acceptors, covers all acid-base reactions in aqueous solution. The strength of acids and bases is related to the percent of their ionization in water. Strong acids and bases are 100 percent ionized, whereas weak acids and bases are less than 5 percent ionized. There are a number of salts that have acidic or basic properties in solution. For example, baking soda, NaHCO3, can be used as an antacid because the bicarbonate ion, HCO3−, is a strong enough conjugate base to combine with H3O+ to give carbonic acid.
HCO3− (aq ) + H3O+ (aq ) → H2CO3 (aq ) + H2O (l ) (27)
The pH scale is a convenient way to represent the acidity or basicity of dilute acid and base solutions. Pure water has a pH of 7; acidic solutions have pH values < 7 and basic solutions have pH values > 7. Each change of one unit of pH is a tenfold change in acidity. Acid-base indicators, such as litmus and phenolphthalein, can be used to measure whether a solution is acidic or basic. A natural "universal indicator," red cabbage extract, can be used to determine the pH within 2 pH units. A buffer contains equal amounts of either a weak acid and its conjugate base or a weak base and its conjugate acid.
see also Arrhenius, Svante; Bases; BrØnsted, Johannes Nicolaus; Chemical Reactions; Lewis, Gilbert N.; Solution Chemistry.
Melvin D. Joesten
Atkins, Peter, and Jones, Loretta (1997). Chemistry: Molecules, Matter, Change, 3rd edition. New York: W. H. Freeman.
Joesten, Melvin D., and Wood, James L. (1996). The World of Chemistry, 2nd edition. Fort Worth, TX: Saunders College.
Moore, John W.; Stanitski, Conrad L.; Wood, James L.; Kotz, John C.; and Joesten, Melvin D. (1998). The Chemical World, 2nd edition. Philadelphia: Saunders.
Shakhashiri, Bassam Z. (1989). Chemical Demonstrations, Vol. 3. Madison: University of Wisconsin Press.
More information available from http://www.visionlearning.com/library/index.htm
"CHEMystery: An Interactive Guide to Chemistry." Available from http://library.thinkquest.org/3659/acidbase/.
"Acid-Base Chemistry." Chemistry: Foundations and Applications. . Encyclopedia.com. (January 16, 2019). https://www.encyclopedia.com/science/news-wires-white-papers-and-books/acid-base-chemistry
"Acid-Base Chemistry." Chemistry: Foundations and Applications. . Retrieved January 16, 2019 from Encyclopedia.com: https://www.encyclopedia.com/science/news-wires-white-papers-and-books/acid-base-chemistry