1. ACIDS and BASES Chapter 18

2. Acids and Bases: An Introduction Acidic solution – contains more hydrogen ions than hydroxide ions. [H + ]>[OH - ] Acidic solution – contains more hydrogen ions than hydroxide ions. [H + ]>[OH - ] Basic solution – contains more hydroxide ions than hydrogen ions. [OH - ]>[H + ] Basic solution – contains more hydroxide ions than hydrogen ions. [OH - ]>[H + ] Arrhenius model – states that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution. Arrhenius model – states that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution. HCl → H + (aq) + Cl - (aq) HCl → H + (aq) + Cl - (aq) NaOH → Na + (aq) + OH - (aq) NaOH → Na + (aq) + OH - (aq) Bronsted-Lowry model – an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. Bronsted-Lowry model – an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. HX(aq) + H 2 O(l) ← → H 3 O + (aq) + X - (aq) HX(aq) + H 2 O(l) ← → H 3 O + (aq) + X - (aq)

3. Acids and Bases Conjugate acid – the species produced when a base accepts a hydrogen ion from an acid. Conjugate acid – the species produced when a base accepts a hydrogen ion from an acid. Conjugate base – the species that results when an acid donates a hydrogen ion to a base. Conjugate base – the species that results when an acid donates a hydrogen ion to a base. Conjugate acid-base pair – consists of two substances related to each other by the donating and accepting of a single hydrogen ion. Conjugate acid-base pair – consists of two substances related to each other by the donating and accepting of a single hydrogen ion. Amphoteric – substances that can act as both acids and bases. (H 2 O) Amphoteric – substances that can act as both acids and bases. (H 2 O) Anhydrides – oxides that become acids or bases by adding the elements contained in water. Anhydrides – oxides that become acids or bases by adding the elements contained in water. CO 2 (g) + H 2 O(l) → H 2 CO 3 (aq) CO 2 (g) + H 2 O(l) → H 2 CO 3 (aq) CaO(s) + H 2 O(l) → Ca 2+ (aq) + 2OH - (l) CaO(s) + H 2 O(l) → Ca 2+ (aq) + 2OH - (l)

4. Strengths of Acids and Bases Strong Acids – acids that ionize completely. Strong Acids – acids that ionize completely. Produce the maximum number of ions, they are good conductors of electricity. HClO 4, HNO 3, HI Produce the maximum number of ions, they are good conductors of electricity. HClO 4, HNO 3, HI HCl(aq) + H 2 O(l) → H 3 O + (aq) + Cl - (aq) HCl(aq) + H 2 O(l) → H 3 O + (aq) + Cl - (aq) Weak Acids – acids that ionizes only partially in dilute aqueous solutions. Weak Acids – acids that ionizes only partially in dilute aqueous solutions. Strong Acids produce a weak conjugate base. Strong Acids produce a weak conjugate base. Weak Acids produce a strong conjugate base. Weak Acids produce a strong conjugate base.

5. Strengths of Acids and Bases Acid ionization constant (K a ) – the value of the equilibrium constant expression for the ionization of a weak acid. Acid ionization constant (K a ) – the value of the equilibrium constant expression for the ionization of a weak acid. Table 19-2 pg.605 Table 19-2 pg.605 Example: Write ionization equations and acid ionization constant expressions for the following acids. Example: Write ionization equations and acid ionization constant expressions for the following acids. HClO 2 & HNO 2 HClO 2 & HNO 2 HClO 2 (aq) + H 2 O(l) ← → H 3 O + (aq) + ClO - (aq) HClO 2 (aq) + H 2 O(l) ← → H 3 O + (aq) + ClO - (aq) K a =[H 3 O + ][ClO 2 - ]/[HClO 2 ] K a =[H 3 O + ][ClO 2 - ]/[HClO 2 ] HNO 2 (aq) + H 2 O(l) ← → H 3 O + (aq) + NO 2 - (aq) HNO 2 (aq) + H 2 O(l) ← → H 3 O + (aq) + NO 2 - (aq) K a =[H 3 O + ][NO 2 - ]/[HNO 2 ] K a =[H 3 O + ][NO 2 - ]/[HNO 2 ]

6. Strengths of Acids and Bases Strong Bases – dissociate entirely into metal ions and hydroxide ions. Strong Bases – dissociate entirely into metal ions and hydroxide ions. Examples: metallic hydroxides, sodium hydroxide Examples: metallic hydroxides, sodium hydroxide NaOH(s) → Na + (aq) + OH - (aq) NaOH(s) → Na + (aq) + OH - (aq) Weak Bases – ionizes only partially in dilute aqueous solution to form the conjugate acid of the base and hydroxide ion. Weak Bases – ionizes only partially in dilute aqueous solution to form the conjugate acid of the base and hydroxide ion. CH 3 NH 2 (aq) + H 2 O(l) ← → CH 3 NH 3 + (aq) + OH - (aq) CH 3 NH 2 (aq) + H 2 O(l) ← → CH 3 NH 3 + (aq) + OH - (aq) Strong bases produce a weak conjugate acid Strong bases produce a weak conjugate acid Weak bases produce a strong conjugate acid Weak bases produce a strong conjugate acid

7. Strengths of Acids and Bases Base ionization constant (K b )- the value of the equilibrium constant expression for the ionization of a base. Base ionization constant (K b )- the value of the equilibrium constant expression for the ionization of a base. Table 19-4 pg.607 Table 19-4 pg.607 Example: Write the ionization equations and base ionization constant expressions for the following bases. Example: Write the ionization equations and base ionization constant expressions for the following bases. Hexylamine (C 6 H 13 NH 2 ) & Propylamine (C 3 H 7 NH 2 ) Hexylamine (C 6 H 13 NH 2 ) & Propylamine (C 3 H 7 NH 2 ) C 6 H 13 NH 2 (aq) + H 2 O(l) ← → C 6 H 13 NH 3 + (aq) + OH - (aq) C 6 H 13 NH 2 (aq) + H 2 O(l) ← → C 6 H 13 NH 3 + (aq) + OH - (aq) K b =[C 6 H 13 NH 3 + ][OH - ]/[C 6 H 13 NH 3 ] K b =[C 6 H 13 NH 3 + ][OH - ]/[C 6 H 13 NH 3 ] C 3 H 7 NH 2 (aq) + H 2 O(l) ← → C 3 H 7 NH 3 + (aq) + OH - (aq) C 3 H 7 NH 2 (aq) + H 2 O(l) ← → C 3 H 7 NH 3 + (aq) + OH - (aq) K b =[C 3 H 7 NH 3 + ][OH - ]/[C 3 H 7 NH 2 ] K b =[C 3 H 7 NH 3 + ][OH - ]/[C 3 H 7 NH 2 ]

8. What is pH? Concentration of pure water. Concentration of pure water. K eq [H 2 O] = K w = [H + ][OH - ] = (1.0x10 -7 )(1.0x10 -7 ) K eq [H 2 O] = K w = [H + ][OH - ] = (1.0x10 -7 )(1.0x10 -7 ) K w =1.0x10 -14 K w =1.0x10 -14 Ion product constant for water – the value of the equilibrium constant expression for the self-ionization of water. Ion product constant for water – the value of the equilibrium constant expression for the self-ionization of water. Example: Calculate [H + ] or [OH - ]. Is the solution acidic, basic, or neutral? Example: Calculate [H + ] or [OH - ]. Is the solution acidic, basic, or neutral? [H + ] = 1.0x10-13M [H + ] = 1.0x10-13M [OH - ] = 1.0x10-7M [OH - ] = 1.0x10-7M

9. What is pH? pH – the negative logarithm of the hydrogen ion concentration. pH – the negative logarithm of the hydrogen ion concentration. pH = -log[H + ] pH = -log[H + ] pOH – negative logarithm of the hydroxide ion concentration. pOH – negative logarithm of the hydroxide ion concentration. pOH = -log[OH - ] pOH = -log[OH - ] pH + pOH = 14.00 pH + pOH = 14.00

10. What is pH? Examples: Calculate the pH or pOH of solutions having the following ion concentrations. Examples: Calculate the pH or pOH of solutions having the following ion concentrations. [H + ] = 1.0x10 -2 M [H + ] = 1.0x10 -2 M [OH - ] = 8.2x10 -6 M [OH - ] = 8.2x10 -6 M [OH-] = 6.5x10-4M [OH-] = 6.5x10-4M [H+] = 0.025M [H+] = 0.025M

11. What is pH? Calculating ion concentrations from pH & pOH Calculating ion concentrations from pH & pOH Antilog(-pH) = [H + ] Antilog(-pH) = [H + ] Antilog(-pOH) = [OH - ] Antilog(-pOH) = [OH - ] Example: Calculate [H + ] and [OH - ]. Example: Calculate [H + ] and [OH - ]. pH = 2.37 pH = 2.37 pH = 11.05 pH = 11.05 Example: Calculate the pH of the following. Example: Calculate the pH of the following. 1.0M HI 1.0M HI 0.050M HNO 3 0.050M HNO 3 1.0M KOH 1.0M KOH

12. What is pH? Example: Calculate the K a for the following acid using the given information. Example: Calculate the K a for the following acid using the given information. 0.0400M solution of HClO 2, pH = 1.80 0.0400M solution of HClO 2, pH = 1.80

13. Neutralization Neutralization reaction – a reaction in which an acid and a base react in aqueous solution to produce a salt and water. Neutralization reaction – a reaction in which an acid and a base react in aqueous solution to produce a salt and water. Double-replacement reaction Double-replacement reaction Salt – an ionic compound made up of a cation from a base and an anion from an acid. Salt – an ionic compound made up of a cation from a base and an anion from an acid. NaOH + HCl → NaCl + H 2 O NaOH + HCl → NaCl + H 2 O base + acid → salt + water base + acid → salt + water

14. Neutralization Example: Write a balanced formula equation for the following acid-base neutralization reaction. Example: Write a balanced formula equation for the following acid-base neutralization reaction. Nitric acid and cesium hydroxide Nitric acid and cesium hydroxide Hydrobromic acid and calcium hydroxide Hydrobromic acid and calcium hydroxide

15. Acid-base Titration Titration – a method for determining the concentration of a solution by reacting a known volume of the solution with a solution of know concentration. Titration – a method for determining the concentration of a solution by reacting a known volume of the solution with a solution of know concentration. Acid-base indicators – chemical dyes whose colors are affected by acidic and basic solutions. Acid-base indicators – chemical dyes whose colors are affected by acidic and basic solutions. End point – point at which the indicator used in a titration changes color. End point – point at which the indicator used in a titration changes color.

16. Acid-base reactions Salt hydrolysis – the anions of the dissociated salt accept hydrogen ions from water or the cations of the dissociated salt donate hydrogen ions to water. Salt hydrolysis – the anions of the dissociated salt accept hydrogen ions from water or the cations of the dissociated salt donate hydrogen ions to water. Buffers – solutions that resist changes in pH when limited amounts of acid or base are added. Buffers – solutions that resist changes in pH when limited amounts of acid or base are added. Buffer capacity – amount of acid or base a buffer solution can absorb without a significant change in pH. Buffer capacity – amount of acid or base a buffer solution can absorb without a significant change in pH.