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Chapter 15 Acids, Bases, and Acid–Base Equilibria: 

Chapter 15 Acids, Bases, and Acid–Base Equilibria General Chemistry: An Integrated Approach Hill, Petrucci, 4th Edition Mark P. Heitz State University of New York at Brockport © 2005, Prentice Hall, Inc.

The Brønsted–Lowry Theory: 

The Brønsted–Lowry Theory According to this theory, an acid is a proton donor and a base is a proton acceptor EOS

Conjugate Acids and Bases: 

Conjugate Acids and Bases EOS The conjugate acid of a base is the base plus the attached proton and the conjugate base of an acid is the acid minus the proton

Weak Acids and Bases: 

Weak Acids and Bases For weak acids and bases, equations can be written to describe equilibrium conditions EOS

Acid–Base Equilibria: 

Acid–Base Equilibria For equilibrium constant expressions, Ka is used to represent the acid ionization constant … … and Kb is used to represent the base ionization constant

Strengths of Conjugate Acid–Base Pairs: 

Strengths of Conjugate Acid–Base Pairs EOS

Acid–Base Strength (cont’d): 

Acid–Base Strength (cont’d) Ka values are used to compare the strengths of weak acids;  K,  strength

Strengths of Binary Acids in a Periodic Group: 

Strengths of Binary Acids in a Periodic Group The greater the tendency for the transfer of a proton from HX to H2O, the more the forward reaction is favored and the stronger the acid EOS

Strengths of Binary Acids in a Periodic Group: 

Strengths of Binary Acids in a Periodic Group Bond-dissociation energy is inversely proportional to acid strength. The weaker the bond, the stronger the acid

Strengths of Other Acids: 

Strengths of Other Acids For oxoacids, as the electronegativity of the central atom increases and as the number of terminal oxygen atoms increases, the acid strength also increases EOS

Strengths of Other Acids: 

Strengths of Other Acids Carboxylic acids all have the -COOH group in common; therefore, differences in acid strength must come from differences in the R group attached to the carboxyl group EOS

Strengths of Amines as Bases: 

Strengths of Amines as Bases Aromatic amines are much weaker bases than aliphatic amines EOS p electrons in the benzene ring of an aromatic molecule are delocalized and can involve the N’s lone-pair electrons in the resonance hybrid

Self-Ionization of Water: 

Self-Ionization of Water Water self-ionizes, that is, it creates a small amount of H3O+ and OH–

Self-Ionization of Water: 

Self-Ionization of Water Kw for water is calculated to be 1.0 × 10–14 M (at 25 oC) EOS This equilibrium constant is very important because it applies to all aqueous solutions—acids, bases, salts, and nonelectrolytes—not just to pure water

pH and pOH: 

pH and pOH

Ionization Constant Relationships: 

Ionization Constant Relationships pKw is the negative logarithm of Kw and at 25 oC is equal to 14.00 pKw = pH + pOH = 14.00 Kw = [H+][OH–] = 1 × 10–14

Acid–Base Equilibrium Calculations: 

Acid–Base Equilibrium Calculations These calculations are similar to the equilibrium calculations performed in Chapter 14

Some Ionization Constants: 

Some Ionization Constants EOS

Polyprotic Acids: 

Polyprotic Acids A monoprotic acid has one ionizable H atom per molecule e.g., HCl, hydrochloric acid

Salts from Acids and Bases: 

Salts from Acids and Bases Salts of strong acids and strong bases form neutral solutions: e.g., NaCl Salts of weak acids and strong bases form basic solutions: e.g., NaF Salts of strong acids and weak bases form acidic solutions: e.g., NH4Cl

The Common Ion Effect: 

The Common Ion Effect If one solution contains a weak acid and another contains that acid and its conjugate base as a second solute, the two solutions have different pH values The conjugate base is referred to as a common ion because it is found in both the weak acid and the anion

Common Ion Effect Illustrated: 

Common Ion Effect Illustrated EOS

Buffer Solutions: 

Buffer Solutions A buffer solution is a solution that changes pH only slightly when small amounts of a strong acid or a strong base are added A buffer is prepared by mixing a weak acid with its salt (conjugate base) or by mixing a weak base with its salt (conjugate acid) in aqueous solution

Buffer Solutions: 

Buffer Solutions The acid component of the buffer can neutralize small added amounts of OH–, and the basic component can neutralize small added amounts of H3O+

Buffering Action: 

Buffering Action EOS

Buffering Action: 

Buffering Action EOS

Buffer Solutions Equation: 

Buffer Solutions Equation The Henderson–Hasselbalch equation is used to calculate pH in a buffer solution as follows:

Buffer Capacity and Buffer Range: 

Buffer Capacity and Buffer Range There is a limit to the capacity of a buffer solution to neutralize added acid or base In general, the more concentrated the buffer components in a solution, the more added acid or base the solution can neutralize

Acid–Base Indicators: 

Acid–Base Indicators

Acid–Base Indicators: 

Acid–Base Indicators

Common Indicators: 

Common Indicators EOS

Neutralization Reactions: 

Neutralization Reactions At the equivalence point in a titration, the acid and base have been brought together in exact stoichiometric proportions mol acid = mol base

Titration Curves: 

Titration Curves

Titration Curve for Weak Acid–Strong Base: 

Titration Curve for Weak Acid–Strong Base Similar features to strong acid/base curve

Titration Curve for Weak Acid–Strong Base: 

Titration Curve for Weak Acid–Strong Base

Lewis Acids and Bases: 

Lewis Acids and Bases There are reactions in nonaqueous solvents, in the gaseous state, and even in the solid state that can be considered acid–base reactions in which Brønsted–Lowry theory is not adequate to explain.

Summary of Concepts: 

Summary of Concepts In the Brønsted–Lowry theory an acid is a proton donor and a base is a proton acceptor If an acid is strong, its conjugate base is weak; and if a base is strong, its conjugate acid is weak Water is amphiprotic: it can be either an acid or a base. It undergoes limited self-ionization producing H3O+ and OH–

Summary (cont’d): 

Summary (cont’d) In aqueous solutions at 25 oC, pH + pOH = 14.00 pH = –log[H3O+] pOH = –log[OH–] pKw = –logKw Hydrolysis reactions cause certain salt solutions to be either acidic or basic A strong electrolyte that produces an ion common to the ionization equilibrium of a weak acid or a weak base suppresses the ionization of the weak electrolyte

Summary (cont’d): 

Summary (cont’d) In Lewis acid–base theory, a Lewis acid accepts an electron pair and a Lewis base donates an electron pair EOS