acids and bases


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acids and bases,

two related classes of chemicals; the members of each class have a number of common properties when dissolved in a solvent, usually water.

Properties

Acids in water solutions exhibit the following common properties: they taste sour; turn litmuslitmus,
organic dye usually used in the laboratory as an indicator of acidity or alkalinity (see acids and bases). Naturally pink in color, it turns blue in alkali solutions and red in acids. Commonly, paper is treated with the coloring matter to form so-called litmus paper.
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 paper red; and react with certain metals, such as zinc, to yield hydrogen gas. Bases in water solutions exhibit these common properties: they taste bitter; turn litmus paper blue; and feel slippery. When a water solution of acid is mixed with a water solution of base, water and a saltsalt,
chemical compound (other than water) formed by a chemical reaction between an acid and a base (see acids and bases). Characteristics and Classification of Salts

The most familiar salt is sodium chloride, the principal component of common table salt.
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 are formed; this process, called neutralizationneutralization,
chemical reaction, according to the Arrhenius theory of acids and bases, in which a water solution of acid is mixed with a water solution of base to form a salt and water; this reaction is complete only if the resulting solution has neither acidic nor basic
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, is complete only if the resulting solution has neither acidic nor basic properties.

Classification

Acids and bases can be classified as organic or inorganic. Some of the more common organic acids are: citric acidcitric acid
or 2-hydroxy-1,2,3-propanetricarboxylic acid,
HO2CCH2C(OH)(CO2H)CH2CO2H, an organic carboxylic acid containing three carboxyl groups; it is a solid at room temperature, melts at 153°C;, and
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, carbonic acidcarbonic acid,
H2CO3, a weak dibasic acid (see acids and bases) formed when carbon dioxide dissolves in water; it exists only in solution. Carbonic acid forms carbonate and bicarbonate (or acid carbonate) salts (see carbonate) by reaction with bases.
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, hydrogen cyanidehydrogen cyanide,
HCN, colorless, volatile, and extremely poisonous chemical compound whose vapors have a bitter almond odor. It melts at −14°C; and boils at 26°C;. It is miscible in all proportions with water or ethanol and is soluble in ether.
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, salicylic acid, lactic acidlactic acid,
CH3CHOHCO2H, a colorless liquid organic acid. It is miscible with water or ethanol. Lactic acid is a fermentation product of lactose (milk sugar); it is present in sour milk, koumiss, leban, yogurt, and cottage cheese.
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, and tartaric acidtartaric acid,
HO2CCHOHCHOHCO2H, white crystalline dicarboxylic acid. It occurs as three distinct isomers, the dextro-, levo-, and meso- forms.
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. Some examples of organic bases are: pyridinepyridine
or azine
, C5H5N, colorless, flammable, toxic liquid with a putrid odor. It melts at −42°C; and boils at 115.5°C;. Chemically, it is a heterocyclic aromatic tertiary amine (see under amino group).
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 and ethylamine. Some of the common inorganic acids are: hydrogen sulfidehydrogen sulfide,
chemical compound, H2S, a colorless, extremely poisonous gas that has a very disagreeable odor, much like that of rotten eggs. It is slightly soluble in water and is soluble in carbon disulfide.
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, phosphoric acidphosphoric acid,
any one of three chemical compounds made up of phosphorus, oxygen, and hydrogen (see acids and bases). The most common, orthophosphoric acid, H3PO4, is usually simply called phosphoric acid.
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, hydrogen chloridehydrogen chloride,
chemical compound, HCl, a colorless, poisonous gas with an unpleasant, acrid odor. It is very soluble in water and readily soluble in alcohol and ether. It fumes in moist air. It is not flammable, and the liquid is a poor conductor of electricity.
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, and sulfuric acidsulfuric acid,
chemical compound, H2SO4, colorless, odorless, extremely corrosive, oily liquid. It is sometimes called oil of vitriol. Concentrated Sulfuric Acid

When heated, the pure 100% acid loses sulfur trioxide gas, SO3
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. Some common inorganic bases are: sodium hydroxidesodium hydroxide,
chemical compound, NaOH, a white crystalline substance that readily absorbs carbon dioxide and moisture from the air. It is very soluble in water, alcohol, and glycerin. It is a caustic and a strong base (see acids and bases).
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, sodium carbonatesodium carbonate,
chemical compound, Na2CO3, soluble in water and very slightly soluble in alcohol. Pure sodium carbonate is a white, odorless powder that absorbs moisture from the air, has an alkaline taste, and forms a strongly alkaline water solution.
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, sodium bicarbonatesodium bicarbonate
or sodium hydrogen carbonate,
chemical compound, NaHCO3, a white crystalline or granular powder, commonly known as bicarbonate of soda or baking soda. It is soluble in water and very slightly soluble in alcohol.
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, calcium hydroxidecalcium hydroxide,
Ca(OH)2, colorless crystal or white powder. It is prepared by reacting calcium oxide (lime) with water, a process called slaking, and is also known as hydrated lime or slaked lime. When heated above 580°C; it dehydrates, forming the oxide.
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, and calcium carbonatecalcium carbonate,
CaCO3, white chemical compound that is the most common nonsiliceous mineral. It occurs in two crystal forms: calcite, which is hexagonal, and aragonite, which is rhombohedral.
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.

Acids, such as hydrochloric acid, and bases, such as potassium hydroxide, that have a great tendency to dissociate in water are completely ionized in solution; they are called strong acids or strong bases. Acids, such as acetic acid, and bases, such as ammonia, that are reluctant to dissociate in water are only partially ionized in solution; they are called weak acids or weak bases. Strong acids in solution produce a high concentration of hydrogen ions, and strong bases in solution produce a high concentration of hydroxide ions and a correspondingly low concentration of hydrogen ions. The hydrogen ion concentration is often expressed in terms of its negative logarithm, or pH. Strong acids and strong bases make very good electrolytes (see electrolysiselectrolysis
, passage of an electric current through a conducting solution or molten salt that is decomposed in the process. The Electrolytic Process

The electrolytic process requires that an electrolyte, an ionized solution or molten metallic salt, complete an
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), i.e., their solutions readily conduct electricity. Weak acids and weak bases make poor electrolytes.

See bufferbuffer,
solution that can keep its relative acidity or alkalinity constant, i.e., keep its pH constant, despite the addition of strong acids or strong bases. Buffer solutions are frequently solutions that contain either a weak acid and one of its salts or a weak base and
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; catalystcatalyst,
substance that can cause a change in the rate of a chemical reaction without itself being consumed in the reaction; the changing of the reaction rate by use of a catalyst is called catalysis.
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; indicators, acid-baseindicators, acid-base,
organic compounds that, in aqueous solution, exhibit color changes indicative of the acidity or basicity of the solution. Common indicators include p-nitrophenol, which is colorless from pH 1 to 5 and yellow from p
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; titrationtitration
, gradual addition of an acidic solution to a basic solution or vice versa (see acids and bases); titrations are used to determine the concentration of acids or bases in solution.
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.

Acid-Base Theories

There are three theories that identify a singular characteristic which defines an acid and a base: the Arrhenius theory, for which the Swedish chemist Svante Arrhenius was awarded the 1903 Nobel Prize in chemistry; the Brönsted-Lowry, or proton donor, theory, advanced in 1923; and the Lewis, or electron-pair, theory, which was also presented in 1923. Each of the three theories has its own advantages and disadvantages; each is useful under certain conditions.

The Arrhenius Theory

When an acid or base dissolves in water, a certain percentage of the acid or base particles will break up, or dissociate (see dissociationdissociation,
in chemistry, separation of a substance into atoms or ions. Thermal dissociation occurs at high temperatures. For example, hydrogen molecules (H2
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), into oppositely charged ions. The Arrhenius theory defines an acid as a compound that can dissociate in water to yield hydrogen ions, H+, and a base as a compound that can dissociate in water to yield hydroxide ions, OH . For example, hydrochloric acid, HCl, dissociates in water to yield the required hydrogen ions, H+, and also chloride ions, Cl . The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH, and also sodium ions, Na+.

The Brönsted-Lowry Theory

Some substances act as acids or bases when they are dissolved in solvents other than water, such as liquid ammonia. The Brönsted-Lowry theory, named for the Danish chemist Johannes Brönsted and the British chemist Thomas Lowry, provides a more general definition of acids and bases that can be used to deal both with solutions that contain no water and solutions that contain water. It defines an acid as a proton donor and a base as a proton acceptor. In the Brönsted-Lowry theory, water, H2O, can be considered an acid or a base since it can lose a proton to form a hydroxide ion, OH, or accept a proton to form a hydronium ion, H3O+ (see amphoterismamphoterism
, in chemistry, the property of certain substances of acting either as acids or as bases depending on the reaction in which they are involved. Many hydroxide compounds are amphoteric.
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). When an acid loses a proton, the remaining species can be a proton acceptor and is called the conjugate base of the acid. Similarly when a base accepts a proton, the resulting species can be a proton donor and is called the conjugate acid of that base. For example, when a water molecule loses a proton to form a hydroxide ion, the hydroxide ion can be considered the conjugate base of the acid, water. When a water molecule accepts a proton to form a hydronium ion, the hydronium ion can be considered the conjugate acid of the base, water.

The Lewis Theory

Another theory that provides a very broad definition of acids and bases has been put forth by the American chemist Gilbert Lewis. The Lewis theory defines an acid as a compound that can accept a pair of electrons and a base as a compound that can donate a pair of electrons. Boron trifluoride, BF3, can be considered a Lewis acid and ethyl alcohol can be considered a Lewis base.

Acids and Bases

 

classes of chemical compounds.

Generally substances are called acids if they contain hydrogen (HCl, HNO3, H2SO4, CH3COOH) and dissociate in water to form H+ ions (more accurately, hydronium ions, H3O+). The presence of these ions determines the sharp taste characteristic of aqueous acid solutions and their ability to change the color of chemical indicators. Acids are classified, by the number of replaceable protons, as monobasic (for example, nitric acid, HNO3; hydrochloric acid, HCl; acetic acid, CH3COOH), dibasic (sulfuric acid, H2SO4; carbonic acid, H2CO3), and tribasic (orthophosphoric acid, H3PO4). The more hydronium ions present in the aqueous acid solution, that is, the higher the degree of dissociation, the stronger the acid. Acids that undergo complete dissociation in dilute solutions are called strong acids. Weak acids have an ionization constant (characterizing the degree of dissociation of the acid in solution, for example, at 25°C) of less than 10–5 (acetic acid, 1.8 × 10–5; prussic acid 7.9 × 10–10). The dissociation of polybasic acids occurs in several stages, each of which is characterized by a particular ionization constant. For example, the ionization constant for the dissociation of H3PO4 into H+ and H2PO4 is 7 × 10–3; for H2PO4 into H+ and HPO42–, 8 × 10–8; and for HPO42– into H+ and PO43–, 4.8 × 10–13.

Generally substances are called bases if they contain the hydroxyl group OH [KOH, NaOH, Ca(OH)2] and undergo dissociation in aqueous solution to form hydroxyl OH ions. The majority of bases are insoluble in water. Water-soluble bases are known as alkalis. The presence of OH ions explains the caustic taste characteristic of alkaline solutions and their ability to change the color of chemical indicators. Bases containing one, two, or three hydroxyl groups are referred to as monoacidic, diacidic, and triacidic, respectively. Bases that do not undergo complete dissociation when dissolved in water are called weak bases. Examples of strong bases are potassium hydroxide, KOH; sodium hydroxide, NaOH; and barium hydroxide, Ba(OH)2.

The concept of acids and bases emerged at the very beginning of the study of chemistry. In 1778 the French chemist A. L. Lavoisier attempted to explain the characteristic features of acids by relative oxygen content. This concept proved unjustified when it became apparent that many oxygen-containing substances (oxides of metals, alkalies, salts) do not exhibit acidic properties and that a number of typical acids (hydrochloric, prussic, hydrofluoric) do not contain oxygen at all. This was demonstrated by the English scientist H. Davy in 1810 and by the French scientist J. L. Gay-Lussac in 1814. The Swedish chemist J. J. Berzelius proposed (1812–19) that acidic and basic properties were due to the electrical character of the oxides—that is, he regarded the electronegative oxides of nonmetals (and certain metals, such as chromium and manganese) as acids and electropositive metal oxides as bases. In 1814, Davy suggested hydrogen as the carrier of acidic properties, since it was a constituent of all compounds known at the time to exhibit those properties. The German chemist J. von Liebig refined this concept substantially in 1833 by stating that the acidic properties of a substance are not determined by the total number of hydrogen atoms contained but only by those atoms that can be replaced by a metal to form salts. After the appearance of the electrolytic dissociation theory of the Swedish scientist S. Arrhenius (1884–87), those compounds were called acids that formed hydrogen H+ ions upon dissociation in aqueous solution, and those that dissociated with the detachment of the hydroxyl OH ion were called bases. The development of the solution theory demonstrated that the interaction both of the substances themselves and of their dissociation products with a solvent plays an important role in electrolytic dissociation. It was also found that the H+ ion cannot exist in a free state in solution. Because of its very high charge density the H+ ion combines stably with the molecules of the solvent (solvation) and actually exists in the form of a solvated ion; in aqueous solutions the H+ actually exists as hydronium ion, which is also a carrier of acidic properties.

A definition of the concepts of acid and base on the basis of the electrolytic dissociation theory is often fully sufficient for practical purposes. However, as has long been established, many compounds exhibiting typically acidic and basic properties contain neither hydrogen nor the hydroxyl group. Furthermore, the same substance frequently behaves as an acid in certain reactions and as a base in others. The ability of a substance to react as either acid or base does not therefore appear to be an absolute property of the substance but is expressed in specific chemical reactions related to the acid-base class. In these reactions, one of the interacting substances acts as an acid in relation to the other substance, which acts as a base. Therefore, the ability of a substance to behave as either acid or base is a functional characteristic. Numerous attempts have been made to develop a unified theory that would make it possible, taking into account the conditions mentioned, unequivocally to call a given substance acid or base. However, no single criterion has yet been found.

The two most popular concepts are those developed by the Danish physical chemist J. N. Brønsted and by the American physical chemist G. N. Lewis (1923). Brønsted classifies as acids hydrogen-containing substances that give up positive hydrogen ions, that is, protons (proton, or Brønsted, acids), and as bases, substances that accept protons. According to Brønsted, both neutral molecules and ions can fulfill the functions of both acids and bases. A chemical reaction involving proton transfer from acid to base (AH + B ⇄ A + BH, where AH is an acid and B is a base) is called an acid-base, or protolytic, reaction. Since protolytic reactions are reversible (proton transfer takes place in both the forward reaction and in the reverse reaction), the products of the forward reaction also fulfill the functions of acid and base (conjugate acids and bases) with respect to each other—that is, BH = acid, and A = base. For example, in the reaction H2SO4 + H2O ⇄ HSO4 + H3O+, the acids are H2SO4 and H3O+, and the bases, HSO4 and H2O. The Brønsted concept offers a precise criterion for relating chemical reactions to the acid-base type. It makes it possible to express the basic characteristics of protolytic equilibriums in quantitative form and to arrange hydrogen-containing substances in a series of increasing ability to give up a proton—according to acidity.

These merits of the protolytic equilibrium theory made for its predicted strength and ensured the broad application of Brønsted’s concepts in practical chemistry. At the same time, these concepts are inherently limited, in that linking the acidic properties of a substance to the presence in it of hydrogen still excludes the large number of acidic substances that contain no hydrogen. Electron-unsaturated compounds (for example, boron, aluminum, and tin halides) and certain metal oxides are among this group, known as aprotonic, or Lewis, acids. According to Lewis’ concept, which to an extent makes up for the above deficiency, an acid is a substance that accepts an electron pair in a chemical reaction and a base is a substance that gives up an electron pair. As a result, the electron unsaturation of an acid molecule is filled by electrons from the base. In addition, a new compound (salt) is formed with a stable electron shell (in particular, an octet) and a donor-acceptor bond. For example,

where BF3 is the acid and NH3 the base.

According to Lewis, an important feature of acid-base reactions is the collectivization of the electron pair of the base. This distinguishes acid-base reactions from oxidation-reduction reactions, which involve the complete removal (one or more at a time) of electrons from the molecules of the reducing agent by the molecules of the oxidizer-carrier; no collectivized orbits result from the process. Unlike Brønsted, Lewis associates acid-base properties not with the presence of specific chemical elements (hydrogen, in particular) but with the structure of the outer electron shells of the atoms alone. At the same time, there is an interrelation between the two concepts, in that a strong affinity for an electron pair is characteristic of both the H+ ion and the Lewis acid. There are certain other concepts about acids and bases but they are not widely accepted.

Both the Brønsted theory and the Lewis theory have found wide practical application. A change in acidity or basicity of a medium is often used to increase the reaction rate and to alter the reaction mechanism; this forms the basis of acid-base catalysis and is used widely in the chemical industry. It should be especially noted that Brønsted and Lewis acids in many cases exhibit similar catalytic activity. Acid-base processes have found wide application in the chemical industry (neutralization, hydrolysis, metal etching). Many acids (sulfuric, nitric, hydrochloric, orthophosphoric) and alkalies (caustic potash, caustic soda) are important both as starting materials and as products in the major branches of the chemical industry.

Acids and bases fulfill diverse structural and dynamic functions in living organisms by taking part in many biological processes. As a rule, these processes are highly sensitive to the acidity or basicity of the medium. The directional effect produced by acids and bases has found application in medicine. For example, weak solutions of hydrochloric acid are used to in-crease gastric secretion, and weak boric acid solution is used as a disinfectant and astringent in mouthwashes. On the other hand, the penetration of concentrated acids and bases into the organism may cause serious burns to internal organs, reduction in cardiac activity, and other damage, frequently resulting in death.

REFERENCE

Luder, W., and S. Zuffante. Elektronnaia teoriia kislot i osnovanii. Moscow, 1950. (Translated from English.)
Usanovich, M. I. Chto takoe kisloty i osnovaniia. Alma-Ata, 1953.
Pauling, L. Obshchaia khimiia. Moscow, 1974. (Translated from English.)
Kratkaia khimicheskaia entsiklopediia, vol. 2. Moscow, 1963.

IA. M. VARSHAVSKII