Chemistry 2e

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The preceding ionization reactions suggest that water may function as both a base (as in its reaction with hydrogen fluoride) and an acid (as in its reaction with ammonia). Species capable of either donating or accepting protons are called amphiprotic, or more generally, amphoteric, a term that may be used for acids and bases per definitions other than the Brønsted-Lowry one. The equations below show the two possible acid-base reactions for two amphiprotic species, bicarbonate ion and water:

Referencing norms: AmphiproticDonatesAndAccepts, NotAmphiproticWithoutAcceptance, NotAmphiproticWithoutDonation

en · official · sha256:6ca53208160e3936a629cdf1430f0d71eb53ec89761f87172495a9b7fc054d94

The acid-base reaction class has been studied for quite some time. In 1680, Robert Boyle reported traits of acid solutions that included their ability to dissolve many substances, to change the colors of certain natural dyes, and to lose these traits after coming in contact with alkali (base) solutions. In the eighteenth century, it was recognized that acids have a sour taste, react with limestone to liberate a gaseous substance (now known to be CO2), and interact with alkalis to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac concluded that acids are substances that can neutralize bases and that these two classes of substances can be defined only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Svante Arrhenius defined an acid as a compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base as a compound that dissolves in water to yield hydroxide anions.

Referencing norms: ArrheniusAcidYieldsHydrogenCation, ArrheniusBaseYieldsHydroxide, ArrheniusReading, ArrheniusReading/AcidIsArrheniusAcid, ArrheniusReading/BaseIsArrheniusBase, ArrheniusReading/NotAcidUnderArrhenius, ArrheniusReading/NotBaseUnderArrhenius, NotArrheniusAcidWhenInsoluble, NotArrheniusAcidWithoutHydrogenCation, NotArrheniusBaseWithoutHydroxide

en · official · sha256:b8b24bbfa29d35b581ced9fc56d6f3a4d7f84075db162e99ad1dc3139bf23c83

The process in which like molecules react to yield ions is called autoionization. Liquid water undergoes autoionization to a very slight extent; at 25 °C, approximately two out of every billion water molecules are ionized. The extent of the water autoionization process is reflected in the value of its equilibrium constant, the ion-product constant for water, Kw: H2O(l)+H2O(l)⇌H3O+(aq)+OH−(aq)Kw=[H3O+][OH−] The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, Kw has a value of 1.0 × 10−14. The process is endothermic, and so the extent of ionization and the resulting concentrations of hydronium ion and hydroxide ion increase with temperature. For example, at 100 °C, the value for Kw is about 5.6 × 10−13, roughly 50 times larger than the value at 25 °C.

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Johannes Brønsted and Thomas Lowry proposed a more general description in 1923 in which acids and bases were defined in terms of the transfer of hydrogen ions, H+. (Note that these hydrogen ions are often referred to simply as protons, since that subatomic particle is the only component of cations derived from the most abundant hydrogen isotope, 1H.) A compound that donates a proton to another compound is called a Brønsted-Lowry acid, and a compound that accepts a proton is called a Brønsted-Lowry base. An acid-base reaction is, thus, the transfer of a proton from a donor (acid) to an acceptor (base).

Referencing norms: ArrheniusAcidIsBronstedAcid, BronstedAcidDonatesProton, BronstedBaseAcceptsProton, BronstedReading, BronstedReading/AcidIsBronstedAcid, BronstedReading/BaseIsBronstedBase, BronstedReading/NotAcidUnderBronsted, BronstedReading/NotBaseUnderBronsted, NotBronstedAcidWithoutDonation, NotBronstedBaseWithoutAcceptance

en · official · sha256:844db56ee19a75002ad1721e93b34023bfea48902c37c1a5ad4ced16adec81a7

The concept of conjugate pairs is useful in describing Brønsted-Lowry acid-base reactions (and other reversible reactions, as well). When an acid donates H+, the species that remains is called the conjugate base of the acid because it reacts as a proton acceptor in the reverse reaction. Likewise, when a base accepts H+, it is converted to its conjugate acid. The reaction between water and ammonia illustrates this idea. In the forward direction, water acts as an acid by donating a proton to ammonia and subsequently becoming a hydroxide ion, OH−, the conjugate base of water. The ammonia acts as a base in accepting this proton, becoming an ammonium ion, NH4+, the conjugate acid of ammonia. In the reverse direction, a hydroxide ion acts as a base in accepting a proton from ammonium ion, which acts as an acid.

en · official · sha256:0aeda026afc9f0ca2c67a40ea011aa3f8020ac6dd97c48088629b35114de9bb5

The reaction between a Brønsted-Lowry acid and water is called acid ionization. For example, when hydrogen fluoride dissolves in water and ionizes, protons are transferred from hydrogen fluoride molecules to water molecules, yielding hydronium ions and fluoride ions: Base ionization of a species occurs when it accepts protons from water molecules. In the example below, pyridine molecules, C5NH5, undergo base ionization when dissolved in water, yielding hydroxide and pyridinium ions:

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And so, at this temperature, acidic solutions are those with hydronium ion molarities greater than 1.0 × 10−7 M and hydroxide ion molarities less than 1.0 × 10−7 M (corresponding to pH values less than 7.00 and pOH values greater than 7.00). Basic solutions are those with hydronium ion molarities less than 1.0 × 10−7 M and hydroxide ion molarities greater than 1.0 × 10−7 M (corresponding to pH values greater than 7.00 and pOH values less than 7.00). Since the autoionization constant Kw is temperature dependent, these correlations between pH values and the acidic/neutral/basic adjectives will be different at temperatures other than 25 °C. For example, the “Check Your Learning” exercise accompanying Example 14.1 showed the hydronium molarity of pure water at 80 °C is 4.9 × 10−7 M, which corresponds to pH and pOH values of:

Referencing norms: AcidicByPh, BasicByPh

en · official · sha256:e5929424fcec432caf58e472a4a478fe9423c64f83c41dc3f181c38d5187d418

At 25 °C, the value of Kw is 1.0 × 10−14, and so: 14.00=pH+pOH As was shown in Example 14.1, the hydronium ion molarity in pure water (or any neutral solution) is 1.0 × 10−7 M at 25 °C. The pH and pOH of a neutral solution at this temperature are therefore: pH=−log[H3O+]=−log(1.0×10−7)=7.00 pOH=−log[OH−]=−log(1.0×10−7)=7.00

Referencing norms: NeutralByPh

en · official · sha256:814fddfb4be60ef540fb66d51b4c853c852daa4891519d0af23a39778c552961

In the absence of any leveling effect, the acid strength of binary compounds of hydrogen with nonmetals (A) increases as the H-A bond strength decreases down a group in the periodic table. For group 17, the order of increasing acidity is HF < HCl < HBr < HI. Likewise, for group 16, the order of increasing acid strength is H2O < H2S < H2Se < H2Te.

en · official · sha256:610b5bc73b43dbed7252fb02f20862bca70674333752615d9ff42beb1491a396

The relative strengths of acids may be quantified by measuring their equilibrium constants in aqueous solutions. In solutions of the same concentration, stronger acids ionize to a greater extent, and so yield higher concentrations of hydronium ions than do weaker acids. The equilibrium constant for an acid is called the acid-ionization constant, Ka. For the reaction of an acid HA: HA(aq)+H2O(l)⇌H3O+(aq)+A−(aq), the acid ionization constant is written Ka=[H3O+][A−][HA] where the concentrations are those at equilibrium. Although water is a reactant in the reaction, it is the solvent as well, so we do not include [H2O] in the equation. The larger the Ka of an acid, the larger the concentration of H3O+ and A− relative to the concentration of the nonionized acid, HA, in an equilibrium mixture, and the stronger the acid. An acid is classified as “strong” when it undergoes complete ionization, in which case the concentration of HA is zero and the acid ionization constant is immeasurably large (Ka ≈ ∞). Acids that are partially ionized are called “weak,” and their acid ionization constants may be experimentally measured. A table of ionization constants for weak acids is provided in Appendix H.

Referencing norms: NotStrongerByIonizationConstant, StrongerByIonizationConstant

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To illustrate this idea, three acid ionization equations and Ka values are shown below. The ionization constants increase from first to last of the listed equations, indicating the relative acid strength increases in the order CH3CO2H < HNO2 < HSO4−: CH3CO2H(aq)+H2O(l)⇌H3O+(aq)+CH3CO2−(aq)Ka=1.8×10−5 HNO2(aq)+H2O(l)⇌H3O+(aq)+NO2−(aq)Ka=4.6×10−4 HSO4−(aq)+H2O(aq)⇌H3O+(aq)+SO42−(aq)Ka=1.2×10−2

Referencing norms: assert-h12de2322f6939d94e0dc4f7a66f3acd3b12805088b0df27bac0ccab006e087de, assert-h26def59d923c095ec851fa79af53823bab9f3b7f4711e0db711fd7d2df0e0bb6, assert-hb93ceb9c93622ef800a11ecdf864344badf44b3b34b9db8b21877ebd3a9a5000

en · official · sha256:4bb1e42d27b73d2e2844cac0e8bda9e91dbd22904cb53f7c3aca975f7211d4ef

If all these strong acids are completely ionized in water, why does the column indicate they vary in strength, with nitric acid being the weakest and perchloric acid the strongest? Notice that the sole acid species present in an aqueous solution of any strong acid is H3O+(aq), meaning that hydronium ion is the strongest acid that may exist in water; any stronger acid will react completely with water to generate hydronium ions. This limit on the acid strength of solutes in a solution is called a leveling effect. To measure the differences in acid strength for “strong” acids, the acids must be dissolved in a solvent that is less basic than water. In such solvents, the acids will be “weak,” and so any differences in the extent of their ionization can be determined. For example, the binary hydrogen halides HCl, HBr, and HI are strong acids in water but weak acids in ethanol (strength increasing HCl < HBr < HI).

Referencing norms: AqueousSolventReading, AqueousSolventReading/LevelingInWater, NonAqueousSolventReading, NonAqueousSolventReading/StrongInNamedSolvent, NonAqueousSolventReading/WeakInNamedSolvent

en · official · sha256:8a13e418e0a1f1f5415efbbb277dceeca46622a64e26e4e8da5ee5ecabfb2bb2

The relative strength of an acid or base is the extent to which it ionizes when dissolved in water. If the ionization reaction is essentially complete, the acid or base is termed strong; if relatively little ionization occurs, the acid or base is weak. As will be evident throughout the remainder of this chapter, there are many more weak acids and bases than strong ones. The most common strong acids and bases are listed in Figure 14.6.

Referencing norms: AqueousSolventReading/StrongInWater, AqueousSolventReading/WeakInWater, StrongBronstedAcidIonizesCompletely, WeakBronstedAcidIonizesPartially

en · official · sha256:ef1bbe725bb8ebf6344ce63df42c956bb0c33ab5cbefabea5adaa4fa7cda136c

Salts with Acidic Ions Salts are ionic compounds composed of cations and anions, either of which may be capable of undergoing an acid or base ionization reaction with water. Aqueous salt solutions, therefore, may be acidic, basic, or neutral, depending on the relative acid-base strengths of the salt's constituent ions. For example, dissolving ammonium chloride in water results in its dissociation, as described by the equation

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Thus, dissolving ammonium chloride in water yields a solution of weak acid cations (NH4+) and inert anions (Cl−), resulting in an acidic solution.

Referencing norms: SaltOfStrongAcidAndWeakBaseIsAcidic

en · official · sha256:3e707c11b15dff24f08df6bd27ff3b1711e3dd00fd4c6a1edf866abec9043498

Dissolving sodium acetate in water yields a solution of inert cations (Na+) and weak base anions (CH3CO2−), resulting in a basic solution.

Referencing norms: SaltOfWeakAcidAndStrongBaseIsBasic

en · official · sha256:414e658845003e41b016e644c7dd9af7c59ba183fc7adc3cc0cb04010dba4f73

Salts with Acidic and Basic Ions Some salts are composed of both acidic and basic ions, and so the pH of their solutions will depend on the relative strengths of these two species. Likewise, some salts contain a single ion that is amphiprotic, and so the relative strengths of this ion’s acid and base character will determine its effect on solution pH. For both types of salts, a comparison of the Ka and Kb values allows prediction of the solution’s acid-base status, as illustrated in the following example exercise.

en · official · sha256:521e149c67ab979441ae9372d79accf7374a9cbbc296ace48319d12584dca3a0

Diprotic acids contain two ionizable hydrogen atoms per molecule; ionization of such acids occurs in two steps. The first ionization always takes place to a greater extent than the second ionization. For example, sulfuric acid, a strong acid, ionizes as follows:

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Acids are classified by the number of protons per molecule that they can give up in a reaction. Acids such as HCl, HNO3, and HCN that contain one ionizable hydrogen atom in each molecule are called monoprotic acids. Their reactions with water are:

en · official · sha256:b4ced2b1dc4c500af7de6e7b3141fc5ddf8cf23a529e21aadef4e48ff3150c89

A solution containing appreciable amounts of a weak conjugate acid-base pair is called a buffer solution, or a buffer. Buffer solutions resist a change in pH when small amounts of a strong acid or a strong base are added (Figure 14.14). A solution of acetic acid and sodium acetate (CH3COOH + CH3COONa) is an example of a buffer that consists of a weak acid and its salt. An example of a buffer that consists of a weak base and its salt is a solution of ammonia and ammonium chloride (NH3(aq) + NH4Cl(aq)).

Referencing norms: BufferIsWeakConjugatePair, NotBufferWithoutTheConjugateBase

en · official · sha256:38d4c0568cc0b3f816971c9c647f7e539cf8d7317b3d9259d39c130e336d5009

Reactions involving the formation of coordinate covalent bonds are classified as Lewis acid-base chemistry. The species donating the electron pair that compose the bond is a Lewis base, the species accepting the electron pair is a Lewis acid, and the product of the reaction is a Lewis acid-base adduct. As the two examples above illustrate, Brønsted-Lowry acid-base reactions represent a subcategory of Lewis acid reactions, specifically, those in which the acid species is H+. A few examples involving other Lewis acids and bases are described below.

Referencing norms: BronstedBaseIsLewisBase, LewisAcidAcceptsElectronPair, LewisBaseDonatesElectronPair, NotLewisAcidWithoutAcceptance, NotLewisBaseWithoutDonation

en · official · sha256:0414b37f7c3aee83f0efda7abc8e6c422a75af41f31611f739de815a4ded1824

The boron atom in boron trifluoride, BF3, has only six electrons in its valence shell. Being short of the preferred octet, BF3 is a very good Lewis acid and reacts with many Lewis bases; a fluoride ion is the Lewis base in this reaction, donating one of its lone pairs:

en · official · sha256:b4f8a6039b78709ad1b88ce4e6096f961b3402bf02b9a5eafc560ba6e3c816e2

In 1923, G. N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to donate a pair of electrons and form a coordinate covalent bond.

Referencing norms: LewisReading, LewisReading/AcidIsLewisAcid, LewisReading/BaseIsLewisBase, LewisReading/NotAcidUnderLewis, LewisReading/NotBaseUnderLewis