CHEMISTRY THE CENTRAL SCIENCE
16 ACID–BASE EQUILIBRIA
CHAPTER SUMMARY AND KEY TERMS
SECTION 16.1 Acids and bases were first recognized by the properties of their aqueous solutions. For example, acids turn litmus red, whereas bases turn litmus blue. Arrhenius recognized that the properties of acidic solutions are due to H+(aq) ions and those of basic solutions are due to OH–(aq) ions.
SECTION 16.2 The Brønsted–Lowry concept of acids and bases is more general than the Arrhenius concept and emphasizes the transfer of a proton (H+) from an acid to a base. The H+ ion, which is merely a proton with no surrounding valence electrons, is strongly bound to water. For this reason, the hydronium ion, H3O+(aq), is often used to represent the predominant form of H+ in water instead of the simpler H+(aq).
A Brønsted–Lowry acid is a substance that donates a proton to another substance; a Brønsted–Lowry base is a substance that accepts a proton from another substance. Water is an amphiprotic substance, one that can function as either a Brønsted–Lowry acid or base, depending on the substance with which it reacts.
The conjugate base of a Brønsted–Lowry acid is the species that remains when a proton is removed from the acid. The conjugate acid of a Brønsted–Lowry base is the species formed by adding a proton to the base. Together, an acid and its conjugate base (or a base and its conjugate acid) are called a conjugate acid–base pair.
The acid–base strengths of conjugate acid–base pairs are related: The stronger an acid, the weaker is its conjugate base; the weaker an acid, the stronger is its conjugate base. In every acid–base reaction, the position of the equilibrium favors the transfer of the proton from the stronger acid to the stronger base.
SECTION 16.3 Water ionizes to a slight degree, forming H+(aq) and OH–(aq). The extent of this autoionization is expressed by the ion-product constant for water: Kw = [H+][OH–] = 1.0 × 10–14 (25 °C). This relationship describes both pure water and aqueous solutions. The Kwexpression indicates that the product of [H+] and [OH–] is a constant. Thus, as [H+] increases, [OH–] decreases. Acidic solutions are those that contain more H+(aq) than OH–(aq), whereas basic solutions contain more OH–(aq) than H+(aq).
SECTION 16.4 The concentration of H+(aq) can be expressed in terms of pH: pH = –log[H+]. At 25 °C the pH of a neutral solution is 7.00, whereas the pH of an acidic solution is below 7.00, and the pH of a basic solution is above 7.00. The pX notation is also used to represent the negative logarithm of other small quantities, as in pOH and pKw. The pH of a solution can be measured using a pH meter, or it can be estimated using acid–base indicators.
SECTION 16.5 Strong acids are strong electrolytes, ionizing completely in aqueous solution. The common strong acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4. The conjugate bases of strong acids have negligible basicity.
Common strong bases are the ionic hydroxides of the alkali metals and the heavy alkaline earth metals.
SECTION 16.6 Weak acids are weak electrolytes; only some of the molecules exist in solution in ionized form. The extent of ionization is expressed by the acid-dissociation constant, Ka, which is the equilibrium constant for the reaction , which can also be written . The larger the value of Ka, the stronger is the acid. For solutions of the same concentration, a stronger acid also has a larger percent ionization. The concentration of a weak acid and its Ka value can be used to calculate the pH of a solution.
Polyprotic acids, such as H2SO3, have more than one ionizable proton. These acids have acid-dissociation constants that decrease in magnitude in the order Ka1 > Ka2 > Ka3. Because nearly all the H+(aq) in a polyprotic acid solution comes from the first dissociation step, the pH can usually be estimated satisfactorily by considering only Ka1.
SECTION 16.7 Weak bases include NH3, amines, and the anions of weak acids. The extent to which a weak base reacts with water to generate the corresponding conjugate acid and OH– is measured by the base-dissociation constant, Kb. This is the equilibrium constant for the reaction , where B is the base.
SECTION 16.8 The relationship between the strength of an acid and the strength of its conjugate base is expressed quantitatively by the equation Ka× Kb = Kw, where Ka and Kb are dissociation constants for conjugate acid–base pairs.
SECTION 16.9 The acid–base properties of salts can be ascribed to the behavior of their respective cations and anions. The reaction of ions with water, with a resultant change in pH, is called hydrolysis. The cations of the alkali metals and the alkaline earth metals as well as the anions of strong acids, such as Cl–, Br–, I–, and NO3–, do not undergo hydrolysis. They are always spectator ions in acid–base chemistry.
SECTION 16.10 The tendency of a substance to show acidic or basic characteristics in water can be correlated with its chemical structure. Acid character requires the presence of a highly polar H—X bond. Acidity is also favored when the H—X bond is weak and when the X– ion is very stable.
For oxyacids with the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom. For oxyacids with the same central atom, acid strength increases as the number of oxygen atoms attached to the central atom increases. Carboxylic acids, which are organic acids containing the COOH group, are the most important class of organic acids. The presence of delocalized pi bonding in the conjugate base is partially responsible for the acidity of these compounds.
SECTION 16.11 The Lewis concept of acids and bases emphasizes the shared electron pair rather than the proton. A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. The Lewis concept is more general than the Brønsted–Lowry concept because it can apply to cases in which the acid is some substance other than H+.
• Define and identify Arrhenius acids and bases. (Section 16.1)
• Understand the nature of the hydrated proton, represented as either H+(aq) or H3O+(aq). (Section 16.2)
• Define and identify Brønsted–Lowry acids and bases and identify conjugate acid–base pairs. (Section 16.2)
• Relate the strength of an acid to the strength of its conjugate base. (Section 16.2)
• Understand how the equilibrium position of a proton-transfer reaction relates the strengths of the acids and bases involved. (Section 16.3)
• Describe the autoionization of water and understand how [H3O+] and [OH–] are related. (Section 16.3)
• Calculate the pH of a solution given [H3O+] or [OH–]. (Section 16.4)
• Calculate the pH of a strong acid or strong base given its concentration. (Section 16.5)
• Calculate Ka or Kb for a weak acid or weak base given its concentration and the pH of the solution. (Sections 16.6 and 16.7)
• Calculate the pH of a weak acid or weak base or its percent ionization given its concentration and Ka or Kb. (Sections 16.6 and 16.7)
• Calculate Kb for a weak base given Ka of its conjugate acid, and similarly calculate Ka from Kb. (Section 16.8)
• Predict whether an aqueous solution of a salt will be acidic, basic, or neutral. (Section 16.9)
• Predict the relative strength of a series of acids from their molecular structures. (Section 16.10)
• Define and identify Lewis acids and bases. (Section 16.11)
Ion product of water at 25 °C
Definition of pH
Definition of pOH
Relationship between pH and pOH
Acid dissociation constant for a weak acid, HA
Percent ionization of a weak acid
Base-dissociation constant for a weak base, B
Relationship between acid- and base-dissociation constants of a conjugate acid–base pair