Section 18.2 Strengths of Acids and Bases Relate the strength of an acid or base to its degree of ionization. Compare the strength of a weak acid with the strength of its conjugate base. Explain the relationship between the strengths of acids and bases and the values of their ionization constants. In solution, strong acids and bases ionize completely, but weak acids and bases ionize only partially. Section 18-2

Acids can be either strong electrolytes or weak electrolytes. Strengths of Acids Acids can be either strong electrolytes or weak electrolytes. Strong acids (such as HCl) completely break up into their ions: HCl (aq)  H+(aq) + Cl-(aq) Weak acids (such as HC2H3O2) only partially break up into their ions: HC2H3O2  H+ (aq) + C2H3O2-(aq) Weak acids don’t completely break up because they go to equilibrium! Strong/Weak Acids

acid ionization constant, Ka Strengths of Acids The equilibrium constant, Keq, provides a quantitative measure of the degree of ionization of an acid. acid ionization constant, Ka Ka is a measure of how much H3O+ is produced The higher the Ka, the stronger the acid Ka indicates whether products or reactants are favored at equilibrium. Section 18-2

Bases can be either strong electrolytes or weak electrolytes. Strengths of Bases Bases can be either strong electrolytes or weak electrolytes. Strong bases (such as NaOH) completely break up into their ions: NaOH (aq)  Na+(aq) + OH-(aq) Weak bases (such as NH3) only partially break up into their ions: NH3 (aq) + H2 O  NH4+ (aq) + OH-(aq) Weak bases don’t completely break up because they go to equilibrium! Strong/Weak Bases

base ionization constant, Kb Strengths of Bases base ionization constant, Kb Kb is a measure of how much OH- is produced the higher the Kb, the stronger the base Section 18-2

A B C D Section 18.2 Assessment A solution with a small Kb is a ____. A. weak acid B. weak base C. strong acid D. strong base A B C D Section 18-2

A B C D Section 18.2 Assessment Where is the equilibrium point in the ionization equation for a strong acid? A. far right B. far left C. slightly right D. slightly left A B C D Section 18-2

Section 18.3 Hydrogen Ions and pH Explain pH and pOH. Relate pH and pOH to the ion product constant for water. Calculate the pH and pOH of aqueous solutions. pH and pOH are logarithmic scales that express the concentrations of hydrogen ions and hydroxide ions in aqueous solutions. Section 18-3

Ion Product Constant for Water Pure water contains equal concentrations of H+ and OH– ions. The ion production of water, Kw = [H+][OH–]. The ion product constant for water is the value of the equilibrium constant expression for the self-ionization of water. Section 18-3

Ion Product Constant for Water (cont.) With pure water at 398 K, both [H+] and [OH–] are equal to 1.0 × 10–7M. Kw at 298 K = 1.0 × 10–14 Kw and LeChâtelier’s Principle proves [H+] × [OH–] must equal 1.0 × 10–14 at 298 K, and as [H+] goes up, [OH–] must go down. Section 18-3

pH and pOH Concentrations of H+ ions are often small numbers expressed in exponential notation. pH is the negative logarithm of the hydrogen ion concentration of a solution. pH = –log [H+] Section 18-3

The sum of pH and pOH equals 14. pH and pOH (cont.) pOH of a solution is the negative logarithm of the hydroxide ion concentration. pOH = –log [OH–] The sum of pH and pOH equals 14. Section 18-3

A B C D Section 18.3 Assessment In dilute aqueous solution, as [H+] increases: A. pH decreases B. pOH increases C. [OH–] decreases D. all of the above A B C D Section 18-3

A B C D Section 18.3 Assessment What is the pH of a neutral solution such as pure water? A. 0 B. 7 C. 14 D. 1.0 × 10–14 A B C D Section 18-3