Acid–Base Concepts: The Brønsted –Lowry Theory

1. HA(aq) H+(aq) + A–(aq) MOH(aq) M+(aq) + OH–(aq) Acid–Base Concepts: The Brønsted–Lowry Theory Arrhenius Acid: A substance that dissociates in water to produce hydrogen ions, H+. Arrhenius Base: A substance that dissociates in water to produce hydroxide ions, OH–.

2. Acid–Base Concepts: The Brønsted–Lowry Theory Brønsted–Lowry Acid: A substance that can transfer hydrogen ions, H+. In other words, a proton donor. Brønsted–Lowry Base: A substance that can accept hydrogen ions, H+. In other words, a proton acceptor. Conjugate Acid–Base Pairs: Chemical species whose formulas differ only by one hydrogen ion, H+.

3. Acid–Base Concepts: The Brønsted–Lowry Theory Acid-Dissociation Equilibrium

4. Acid–Base Concepts: The Brønsted–Lowry Theory Base-Dissociation Equilibrium

5. HA(aq) + H2O(l) H3O+(aq) + A–(aq) Stronger acid + Stronger base Weaker acid + Weaker base Acid Strength and Base Strength Acid Base Acid Base With equal concentrations of reactants and products, what will be the direction of reaction?

6. Acid Strength and Base Strength Weak Acid: An acid that is only partially dissociated in water and is thus a weak electrolyte.

7. Acid Strength and Base Strength

8. Factors That Affect Acid Strength Bond Strength

9. Factors That Affect Acid Strength Bond Polarity

10. Factors That Affect Acid Strength Oxoacids

11. Factors That Affect Acid Strength Oxoacids

12. 2 H2O(l) H3O+(aq) + OH–(aq) Dissociation of Water Dissociation of Water: Ion-Product Constant for Water: Kw = [H3O+][OH–] at 25 °C: [H3O+] = [OH–] = 1.0 × 10–7 M Kw = (1.0 × 10–7)(1.0 × 10–7) = 1.0 × 10–14

13. 1.0 × 10–14 1.0 × 10–14 Kw [H3O+] [OH–] [H3O+] = = [OH–] Kw [OH–] = = [H3O+] Dissociation of Water Kw = [H3O+][OH–] = 1.0 × 10–14

14. Dissociation of Water

15. [H3O+] = 10 –pH Basic solution: pH > 7 Neutral solution: pH = 7 Acidic solution: pH < 7 The pH Scale pH = –log[H3O+]

16. The pH Scale The hydronium ion concentration for lemon juice is approximately 0.0025 M. What is the pH of lemon juice? pH = –log(0.0025) = 2.60 2 decimal places, is the normal protocol

17. Summary of Acid/Base Equations • pH = - log [H+] • pOH = - log [OH-] • 10-pH = [H+] • 10-pOH

18. 1.0 × 10–14 1.0 × 10–14 [OH–] 0.0019 The pH Scale Calculate the pH of an aqueous ammonia solution that has an OH– concentration of 0.0019 M. [H3O+] = = = 5.3 × 10–12 M pH = –log(5.3 × 10–12) = 11.28

19. [H3O+] = 10 –4.5 The pH Scale Acid rain is a matter of serious concern because most species of fish die in waters having a pH lower than 4.5–5.0. Calculate [H3O+] in a lake that has a pH of 4.5. = 3 × 10–5 M

20. HIn(aq) + H2O(l) H3O+(aq) + In–(aq) Measuring pH Acid–Base Indicator: A substance that changes color in a specific pH range. Indicators exhibit pH-dependent color changes because they are weak acids and have different colors in their acid (HIn) and conjugate base (In–) forms. Color A Color B

21. HNO3(aq) + H2O(l) H3O+(aq) + NO3–(aq) The pH in Solutions of Strong Acids and Strong Bases What is the pH of a 0.025 M solution of HNO3? 100% Since HNO3 is a strong acid, [H3O+] = [HNO3]. pH = –log([H3O+]) = –log(0.025) = 1.60

22. 1.0 × 10–14 1.0 × 10–14 [OH–] 0.025 NaOH(aq) Na+(aq) + OH–(aq) The pH in Solutions of Strong Acids and Strong Bases What is the pH of a 0.025 M solution of NaOH? Since NaOH is a strong base, [OH–] = [NaOH]. [H3O+] = = = 4.0 × 10–13 M pH = –log([H3O+]) = –log(4.0 × 10–13 ) = 12.40

23. HA(aq) + H2O(l) H3O+(aq) + A–(aq) [H3O+][A–] Ka = [HA] Equilibria in Solutions of Weak Acids Acid-Dissociation Constant:

24. Equilibria in Solutions of Weak Acids

25. HF(aq) + H2O(l) H3O+(aq) + F–(aq) Equilibria in Solutions of Weak Acids The pH of 0.250 M HF is 2.036. What are the values of Ka and pKa for hydrofluoric acid? x = [H3O1+] = 10–2.036 = 0.00920 M

26. [H3O+][F–] [H3O+][F–] Ka = [HF] [HF] (0.00920)(0.00920) 0.241 Equilibria in Solutions of Weak Acids [F–] = [H3O+] = 0.00920 M [HF] = 0.250 – x = 0.250 – 0.00920 = 0.241 M Ka = = = 3.51 × 10–4 pKa = –log(Ka) = –log(3.51 × 10–4) = 3.455

27. HCN(aq) + H2O(l) H3O+(aq) + CN–(aq) [H3O+][CN–] Ka = [HCN] Calculating Equilibrium Concentrations in Solutions of Weak Acids Calculate the pH of a 0.10 M HCN solution. At 25 °C, Ka = 1.4 × 10–9.

28. (x)(x) x2 (0.10 – x) 0.10 Calculating Equilibrium Concentrations in Solutions of Weak Acids 4.9 × 10–10 = ≈ x = [H3O+] = 7.0 × 10–6 M pH =–log([H3O+]) = –log(7.0 × 10–6) = 5.15

30. [HA] dissociated [HA] initial Percent Dissociation in Solutions of Weak Acids × 100% Percent dissociation =

31. [BH+][OH–] NH3(aq) + H2O(l) B(aq) + H2O(l) NH4+(aq) + OH–(aq) BH+(aq) + OH–(aq) Kb = [B] [NH4+][OH–] [NH3] Equilibria in Solutions of Weak Bases Base Acid Acid Base Base-Dissociation Constant: Kb =

32. Equilibria in Solutions of Weak Bases

33. NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) [NH4+][OH–] Kb = [NH3] Equilibria in Solutions of Weak Bases Calculate the pH of a 0.40 M NH3 solution. At 25 °C, Kb = 1.8 × 10–5.

34. (x)(x) x2 (0.40 – x) 0.40 1.0 × 10–14 0.0027 Equilibria in Solutions of Weak Bases 1.8 × 10–5 = ≈ x = [OH–] = 0.0027 M [H3O+] = = 3.7 × 10–12 M pH = –log([H3O+]) = –log(3.7 × 10–12) = 11.43

35. NH4+(aq) + H2O(l) NH3(aq) + H2O(l) H3O+(aq) + NH3(aq) NH4+(aq) + OH–(aq) 2H2O(l) H3O+(aq) + OH–(aq) Relation Between Ka and Kb Ka Kb Kw [H3O+][NH3] [NH4+][OH–] Ka × Kb = × = [H3O+][OH–] = Kw [NH4+] [NH3] = (5.6 × 10–10)(1.8 × 10–5) = 1.0 × 10–14

36. Kw Kw Ka Kb Relation Between Ka and Kb Ka × Kb = Kw conjugate acid–base pair Ka = Kb = pKa + pKb = pKw = 14.00

37. Acid–Base Properties of Salts

38. Acid–Base Properties of Salts Salts That Yield Neutral Solutions The following ions do not react appreciably with water to produce either H3O+ or OH– ions: • Cations from strong bases: • Alkali metal cations of group 1A (Li+, Na+, K+) • Alkaline earth metal cations of group 2A (Mg2+, Ca2+, Sr2+, Ba2+), except for Be2+ • Anions from strong monoprotic acids: • Cl–, Br–, I–, NO3–, and CIO4–

39. NH4+(aq) + H2O(l) H3O+(aq) + NH3(aq) Acid–Base Properties of Salts Salts That Yield Acidic Solutions Salts such as NH4Cl that are derived from a weak base (NH3) and a strong acid (HCl) yield acidic solutions. Ammonium ion (NH4+) is the conjugate acid of the weak base ammonia (NH3), while chloride ion (Cl–) is neither acidic nor basic.

40. CN–(aq) + H2O(l) HCN(aq) + OH–(aq) Acid–Base Properties of Salts Salts That Yield Basic Solutions Salts such as NaCN that are derived from a strong base (NaOH) and a weak acid (HCN) yield basic solutions. Cyanide ion (CN–) is the conjugate base of the weak acid hydrocyanic acid (HCN), while sodium ion (Na+) is neither acidic nor basic.

41. Acid–Base Properties of Salts Salts That Contain Acidic Cations and Basic Anions The pH of an ammonium carbonate solution, (NH4)2CO3, depends on the relative acid strength of the cation and the relative base strength of the anion. Is it acidic or basic?

42. CO32–(aq) + H2O(l) NH4+(aq) + H2O(l) HCO3–(aq) + OH–(aq) H3O+(aq) + NH3(aq) Acid–Base Properties of Salts Salts That Contain Acidic Cations and Basic Anions (NH4)2CO3: Ka Kb Three possibilities: • Ka > Kb: The solution will contain an excess of H3O+ ions (pH < 7). • Ka < Kb: The solution will contain an excess of OH– ions (pH > 7). • Ka ≈ Kb: The solution will contain approximately equal concentrations of H3O+ and OH– ions (pH ≈ 7).

43. CO32–(aq) + H2O(l) NH4+(aq) + H2O(l) HCO3–(aq) + OH–(aq) H3O+(aq) + NH3(aq) 1.0 × 10–14 1.0 × 10–14 Kw 5.6 × 10–11 1.8 × 10–5 Kb for NH3 Kw Ka for HCO3– Acid–Base Properties of Salts Salts That Contain Acidic Cations and Basic Anions (NH4)2CO3: Ka Kb Ka for NH4+ = = = 5.6 × 10–10 Kb for CO32– = = = 1.8 × 10–4 Basic, Ka < Kb

44. Acid–Base Properties of Salts - Summary