Acids And Bases

I.

Acids and Bases

INTRODUCTION

Acids and Bases, two classes of chemical compounds that display generally opposite characteristics. Acids taste sour, turn litmus (a pink dye derived from lichens) red, and often react with some metals to produce hydrogen gas. Bases taste bitter, turn litmus blue, and feel slippery. When aqueous (water) solutions of an acid and a base are combined, a neutralization reaction occurs. This reaction is characteristically very rapid and generally produces water and a salt. For example, sulfuric acid and sodium hydroxide, NaOH, yield water and sodium sulfate:

H2SO4 + 2NaOH⇄2H2O + Na2SO4

pH Scale: Some Common Solutions The pH of a solution measures the hydrogen ion concentration in that solution. A small change in pH represents a large change in hydrogen ion concentration. For example, the hydrogen ion concentration of lemon juice (pH of 2.3) is 63 times greater than that of tomato juice (pH of 4.1), and 50,000 times greater than that of water (pH of 7.0).

Common Household Acids and Bases Acids are substances that release hydrogen ions (hydrogen atoms missing their single electron) when they are in a water solution. Bases are substances that accept hydrogen ions in a water solution. Name

Formula Locations

Acids Acetic acid

HC2H3O2 Vinegar (aqueous solution)

Acetylsalicylic acid

HC9H7O4 Aspirin

Ascorbic acid

H2C6H6O6 Vitamin C

Citric acid

H3C6H5O7 Lemon juice, citrus fruits

Hydrochloric acid

HCl

Gastric juices (digestive fluid in stomach)

Sulfuric acid

H2SO4

Batteries

Ammonia

NH3

Household cleaners (aqueous solution)

Calcium hydroxide

Ca(OH)2 Slaked lime (used in mortar for construction)

Magnesium hydroxide

Mg(OH)2 Milk of magnesia (antacid and laxative)

Potassium hydroxide (also called caustic potash)

KOH

Soft soap

Sodium hydroxide

NaOH

Drain and oven cleaners

Bases

II.

EARLY THEORIES

Modern understanding of acids and bases began with the discovery in 1834 by the English physicist Michael Faraday that acids, bases, and salts are electrolytes. That is, when they are dissolved in water, they produce a solution that contains charged particles, or ions, and can conduct an electric current Ionization. In 1884 the Swedish chemist Svante Arrhenius (and later Wilhelm Ostwald, a German chemist) proposed that an acid be defined as a hydrogen-containing compound that, when dissolved in water, produces a concentration of hydrogen ions, or protons, greater than that of pure water. Similarly, Arrhenius proposed that a base be defined as a substance that, when dissolved in water, produces an excess of hydroxyl ions, OH-. The neutralization reaction then becomes: H+ + OH-⇄H2O A number of criticisms of the Arrhenius-Ostwald theory have been made. First, acids are restricted to hydrogen-containing species and bases to hydroxylcontaining species. Second, the theory applies to aqueous solutions exclusively, whereas many acid-base reactions are known to take place in the absence of water.

III.

BRØNSTED-LOWRY THEORY

A more satisfactory theory was proposed in 1923 by the Danish chemist Johannes Brønsted and independently by Thomas Lowry, a British chemist. Their theory states that an acid is a proton (hydrogen ion, H+) donor and a base a proton acceptor. Although the acid must still contain hydrogen, the BrønstedLowry theory does not require an aqueous medium. For example, liquid ammonia, which acts as a base in aqueous solution, can act as an acid in the absence of water by transferring a proton to a base and forming the amide anion (negative ion) NH2-:

NH3 + base⇄NH2- + base + H+

The Brønsted-Lowry definition of acids and bases also explains why a strong acid displaces a weak acid from its compounds (and likewise for strong and weak bases). Here acid-base reactions are viewed as a competition for protons. In terms of a general chemical equation, the reaction of Acid (1) with Base (2)

Acid (1) + Base (2)⇄Acid (2) + Base (1) results in the transfer of a proton from Acid (1) to Base (2). In losing the proton, Acid (1) becomes its conjugate base, Base (1). In gaining a proton, Base (2) becomes its conjugate acid, Acid (2). The equilibrium represented by the equation above may be displaced either to the left or to the right, and the actual reaction will take place in the direction that produces the weaker acid-base pair. For example, hydrogen chloride (HCl) is a strong acid in water because it readily transfers a proton to water to form a hydronium ion:

HCl + H2O⇄H3O+ + Cl-

The equilibrium lies mostly to the right because the conjugate base of HCl, Cl-, is a weak base, and H3O+, the conjugate acid of H2O, is a weak acid. In contrast, hydrogen fluoride, HF, is a weak acid in water because it does not readily transfer a proton to water:

HF + H2O⇄H3O+ + F-

This equilibrium lies mostly to the left because H2O is a weaker base than F-, and because HF is a weaker acid (in water) than H3O+. The Brønsted-Lowry theory also explains why water can be amphoteric, that is, why it can serve as either an acid or a base. Water serves as a base in the presence of an acid that is stronger than water (such as HCl), in other words, an acid that has a greater tendency to dissociate than does water:

HCl + H2O⇄H3O+ + Cl-

Water can also serve as an acid in the presence of a base that is stronger than water (such as ammonia):

IV.

NH3 + H2O⇄NH4+ + OHLEWIS ACID BASE THEORY

In 1923 G. N. Lewis suggested another way of looking at the reaction between H+ and OH- ions. In the Brnsted model, the OH- ion is the active species in this reaction it accepts an H+ ion to form a covalent bond. In the Lewis model, the H+ ion is the active species it accepts a pair of electrons from the OH- ion to form a covalent bond.

In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of electrons. A Lewis acid is therefore any substance, such as the H+ ion,

that can accept a pair of nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such as the OH- ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair donor. One advantage of the Lewis theory is the way it complements the model of oxidationreduction reactions. Oxidation-reduction reactions involve a transfer of electrons from one atom to another, with a net change in the oxidation number of one or more atoms.

The Lewis theory suggests that acids react with bases to share a pair of electrons, with no change in the oxidation numbers of any atoms. Many chemical reactions can be sorted into one or the other of these classes. Either electrons are transferred from one atom to another, or the atoms come together to share a pair of electrons. The principal advantage of the Lewis theory is the way it expands the number of acids and therefore the number of acid-base reactions. In the Lewis theory, an acid is any ion or molecule that can accept a pair of nonbonding valence electrons. In the preceding section, we concluded that Al3+ ions form bonds to six water molecules to give a complex ion.

Al3+(aq) + 6 H2O(l)

Al(H2O)63+(aq)

This is an example of a Lewis acid-base reaction. The Lewis structure of water suggests that this molecule has nonbonding pairs of valence electrons and can therefore act as a Lewis base. The electron configuration of the Al3+ ion suggests that this ion has empty 3s, 3p, and 3d orbitals that can be used to hold pairs of nonbonding electrons donated by neighboring water molecules. Al3+ = [Ne] 3s0 3p0 3d0 Thus, the Al(H2O)63+ ion is formed when an Al3+ ion acting as a Lewis acid picks up six pairs of electrons from neighboring water molecules acting as Lewis bases to give an acid-base complex, or complex ion. The Lewis acid-base theroy explains why BF3 reacts with ammonia. BF3 is a trigonalplanar molecule because electrons can be found in only three places in the valence shell of the boron atom. As a result, the boron atom is sp2 hybridized, which leaves an empty 2pz orbital on the boron atom. BF3 can therefore act as an electron-pair acceptor, or Lewis acid. It can use the empty 2pz orbital to pick up a pair of nonbonding electrons from a Lewis base to form a covalent bond. BF3 therefore reacts with Lewis bases such as NH3 to form acid-base complexes in which all of the atoms have a filled shell of valence electrons, as shown in the figure below.

The Lewis acid-base theory can also be used to explain why nonmetal oxides such as CO2 dissolve in water to form acids, such as carbonic acid H2CO3.

CO2(g) + H2O(l)

H2CO3(aq)

In the course of this reaction, the water molecule acts as an electron-pair donor, or Lewis base. The electron-pair acceptor is the carbon atom in CO2. When the carbon atom picks up a pair of electrons from the water molecule, it no longer needs to form double bonds with both of the other oxygen atoms as shown in the figure below

One of the oxygen atoms in the intermediate formed when water is added to CO2 carries a positive charge; another carries a negative charge. After an H+ ion has been transferred from one of these oxygen atoms to the other, all of the oxygen atoms in the compound are electrically neutral. The net result of the reaction between CO2 and water is therefore carbonic acid, H2CO3.

V.

MEASURING ACID OR BASE STRENGTH

Relative Strengths of Common Acids and Bases The relative strength of acids and bases depends on their tendency to donate or accept hydrogen ions (hydrogen atoms missing their single electron). Strong acids lose hydrogen ions easily, while strong bases accept hydrogen ions easily. The conjugate of a substance is the substance whose chemical formula differs from the

first’s formula by one hydrogen ion, or H+. The conjugate of a strong acid is a weak base, and the conjugate of a strong base is a weak acid. The strength of an acid can be measured by the extent to which an acid transfers a proton to water to produce the hydronium ion, H3O+. Conversely, the strength of a base is indicated by the extent to which the base removes a proton from water. A convenient acid-base scale is calculated from the amount of H3O+ that is formed in water solutions of acids or of OH- formed in water solutions of bases. The former is known as the pH scale and the latter as the pOH scale (pH). The value for pH is equal to the negative logarithm of the hydronium ion concentration—and for pOH, of the hydroxyl ion concentration—in an aqueous solution: pH = -log [H3O+] pOH = -log [OH-] Pure water has a pH of 7.0. When an acid is added, the hydronium ion concentration [H3O+] becomes larger than that in pure water, and the pH becomes less than 7.0, depending on the strength of the acid. The pOH of pure water is also 7.0, and in the presence of a base, the pOH drops to values lower than 7.0. The American chemist Gilbert N. Lewis has offered another theory of acids and bases that has the further advantage of not requiring the acid to contain hydrogen. This theory states that acids are electron-pair acceptors and bases are electron-pair donors. This theory also has the advantages that it works when solvents other than water are involved and it does not require the formation of a salt or of acid-base conjugate pairs. Thus, ammonia is viewed as a base because it can donate an electron pair to the acid boron trifluoride, for example H3N: + BF3⇄H3N-BF3 to form an acid-base association pair.

MORE SOURCES Web Links Acid-base without algebra: graphical treatment of acid-base systems This privately maintained site presents a visual approach to solving acid-base problems; it includes illustrated tutorials and links to relevant Web sites. http://www2.sfu.ca/person/lower/TUTORIALS/AQCHEM/ABG-ind.html Acids and Bases Problem Set The Biology Project at the University of Arizona offers a problem set on the characteristics of acids and bases; explanations and interactive tutorials are included. http://www.biology.arizona.edu/biochemistry/problem_sets/ph/ph.html

Further Reading Chemistry Atkins, Peter Williams. The Periodic Kingdom: A Journey into the Land of the Chemical Elements. Perseus , 1997. An introductory text that inventively serves as a travel guide to the field of chemistry. Hall, Nina. The New Chemistry. Cambridge University Press, 2000. This book highlights the most important developments in chemistry over the past 30 years. Hill, John W., and Doris W. Kolb. Chemistry for Changing Times. 8th ed. Prentice Hall, 1997. Application of the essentials of chemistry to food, drugs, air, etc. Joesten, Melvin D. and others. World of Chemistry. 2nd ed. Saunders, 1997. Essays by prominent chemists describing for lay people or beginners just what their interest are. Keinan, Ehud, and Israel Schechter. Chemistry for the 21st Century. Wiley, 2001. Provides a glimpse into the future of this dynamic science. Malone, Lee J. Basic Concepts of Chemistry. 6th ed. Wiley, 2000. A student-friendly introduction to chemistry. Snyder, Carl H. The Extraordinary Chemistry of Ordinary Things. 3rd ed. Wiley, 1997. Chemistry for those who will not pursue it professionally. Contributed By: Paul L. Gaus Microsoft ® Encarta ® 2006. © 1993-2005 Microsoft Corporation. All rights reserved.