Acids

1. Acids • Bronsted-Lowry Theory: acids donate protons (H+) in a chemical reaction. • Ex. HCl(g) + H2O(l) --> H3O+(aq) + Cl-(aq)

2. Acids • The HCl gas donates a proton to the water molecule, producing the hydronium ion. HCl is considered an acid, water is considered a base.

3. Acids • Bronsted-Lowry Theory: bases accept protons (H+) in a chemical reaction. • HCl and Cl- are considered a conjugate acid/base pair.

4. Acids • Conjugate acid/base pairs: Conjugate base - the particle leftover after the acid donates a proton.

5. Acids • Conjugate acid/base pairs: Conjugate acid - the particle produced after the base accepts the proton.

6. Acid Nomenclature • Binary acids: acids consisting of 2 elements. • Ex. HCl • HI • HBr

7. Acid Nomenclature • Ternary Acids and Bases: acids or bases containing three elements. • Common ternary acid - formed by using H+ and a common polyatomic ion.

8. Acid Nomenclature • Use the polyatomic name and the suffix -ic. • Ex. H2SO4 • HNO3 • HClO3

9. Acid Nomenclature • A ternary acid that is the same as the common acid but with one less oxygen uses the suffix -ous. • Ex. H2SO3 • HClO2

10. Acid Nomenclature • A ternary acid that is the same as the common acid but with two less oxygens uses the prefix -hypo and the suffix -ous. • Ex. HClO

11. Acid Nomenclature • A ternary acid that is the same as the common acid but with one more oxygen uses the prefix -per and the suffix -ic. • Ex. HClO4

12. Acid Nomenclature • Common ternary base - formed by using metal and the hydroxide polyatomic ion. • Ex. NaOH, Mg(OH)2

13. Acid Nomenclature • Organic acids: carboxylic acids, -COOH • Name the chain and add -oic acid.

14. Water Equilibria • Water self ionizes. That is, experiments have shown that a small amount of water becomes H+ and OH-

15. Water Equilibrium • In any aqueous solution, there is an equilibrium between H3O+ (H+) ions and OH- ions. • H2O <--> H+ (aq) + OH- (aq)

16. Water Equilibria • Ion product constant of water: • Kw = Keq[H2O] = [H3O+][OH-] = (1.00 x 10-14)

17. Water Equilibria • What is the hydroxide concentration in a water solution with [H3O+]= 1.5 x 10-6?

18. Water Equilibrium • In pure water: [H+] = [OH-] = 1.0 x 10-7 M; neutral solution.

19. Water Equilibrium • Acidic Solution: [H+] > 1.0 x 10-7 M > [OH-]. • Basic Solution: [H+] < 1.0 x 10-7 M < [OH-].

20. pH Scale • Measures the hydronium concentration in a solution, and thus, its acidity. • pH = -log[H3O+]

21. pH Scale • The lower the number, the more acidic the solution. Numbers range from 0-14 (<7 = acid, 7= neutral, >7= basic)

22. pH Scale • A pH change of 1 = a 10x change in acid concentration. • pOH

23. Water Equilibria Suppose [H+] = 2.4 x 10-6 M; calculate pH. pH = 5.62

24. Water Equilibria Suppose pH = 8.68; calculate [H+]. [H+] = 10-8.68 = 2.1 x 10-9 M.

25. Ionization Constant • The Keq for the ionization of a weak acid or base determines the extent that [H3O+] or [OH-] ions will be produced at equilibrium.

26. Ionization Constant • Ka is the ionization constant of a weak acid. HX + H2O <--> H3O+ + X-. • What is Ka for this reaction?

27. Ionization Constant • Kb is the ionization constant of a weak base. NH3+ H2O <--> NH4+ + OH-. • What is Kb for this reaction?

28. Ionization Constant • Percent ionization = [amount ionized] / [original acid] x 100%

29. Ionization Constant • Weak Acids: molecules • Ex. HF(aq) + H2O <--> H3O+(aq) + F-(aq)

30. Ionization Constant • Weak Acids: cations • Ex. NH4+(aq) + H2O <--> H3O+(aq) + NH3(aq)

31. Ionization Constant • Weak Bases: molecules • Ex. NH3(aq) + H2O <--> OH-(aq) + NH4+(aq)

32. Ionization Constant • Weak Bases: anions derived from weak acids • Ex. F-(aq) + H2O <--> HF(aq) + OH-(aq)

33. Ionization Constant The pH of 0.100 M HC2H3O2 = 2.87; calculate Ka. [H+] = [C2H3O2-] = 1.3 x 10-3 M [HC2H3O2] = 0.100 M - 0.0013 M = 0.099 M. Ka = 1.7 x 10-5

34. Ionization Constant Find [H+] in 0.200 M HC2H3O2 , Ka = 1.8 x 10-5. [H+] = 1.9 x 10-3 M

35. Ionization Constant Note: let x = [H+]. The change in initial acid concentration is often negligible. Assume this to be true…

36. Ionization Constant …If this estimated x is less than 5% of the original concentration, then this answer is considered acceptable. If not then you must make a second approximation.

37. Ionization Constant Find [H+] in 0.100 M HF, Ka = 6.9 x 10-4. [H+] = 8.0 x 10-3 M (you must do a second approximation).

38. Ionization Constant • Ka x Kb = Kw = 1.0 x 10-14 • pH + pOH = 14 • pKa + pKb = 14

39. Ionization Constant • Ka vs Kb

40. Ionization Constant What is the [OH-] and pH of 0.10 M NaF? [OH-] = 1.2 x 10-6 M; pH = 8.08

41. Ionization Constant What is Keq for the reaction between sodium fluoride and phosphoric acid? Phosphoric acid Ka = 7.5x10-3 Hydrofluoric acid Ka = 3.5x10-4 Keq = 21.4

42. Ionization Constant • Common Ion Effect - adding ions that are the same as one of those produced by the ionization of a weak electrolyte …. (more)

43. Ionization Constant • to a solution of the electrolyte suppresses its ionization. Why???

44. Ionization Constant • Le Chatelier’s Principle - you are increasing the concentration of products. The system responds to reestablish equilibrium.

45. Acid/Base Behavior • Consider a compound in the form HOX. If X is very electronegative then the H is given up as a proton and it acts as an acid. If not it acts as a base.

46. Acid/Base Behavior • Ex.NaOH • Nonmetals oxides tend to form acids, metal oxides tend to form bases when dissolved in water. Ex. MgO, CO

47. Acid/Base Behavior • Acidic and Basic Anhydrides: acids and bases that have had water removed. • Ex. Acid anhydride + water --> acid

48. Acid/Base Behavior • Ex. Acid anhydride + water --> acid • SO2 + H2O --> H2SO3

49. Acid/Base Behavior • Ex. basic anhydride + water --> base • Na2O + H2O --> 2 NaOH

50. Salts and Solutions • Salt : an ionic compound that does not consist of H+ or OH- • Ex. KCl, MgO (any ionic compound that is not an Arrehnius acid or base)