Acids and Bases

1. Acids and Bases

2. Acids & Bases The Bronsted-Lowry model defines an acid as a proton donor. A base is a proton acceptor. Note that this definition is based on the transfer of a proton from the acid to the base.

3. NH3(aq) + H3O+(l)

4. Acids & Bases H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) In this reaction, water accepts a proton from HCl. Water is a base, and HCl is an acid. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor B-L base B-L acid

5. Acids & Bases Since this reaction goes to completion (note the one-way arrow), we classify HCl(aq) as a strong acid. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor

6. Acids & Bases Hydrochloric acid dissociates 100%, and exists as a solution of hydronium and chloride ions. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor Although a bottle may be labeled 1.0M HCl, it really contains 1.0 M H3O+(aq) and 1.0 M Cl1-(aq).

7. Acids & Bases Hydrochloric acid dissociates 100%, and exists as a solution of hydronium and chloride ions. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor There is no reverse reaction because chloride has no tendency to accept a proton to form HCl.

8. Acids & Bases There are only a few common strong acids. They are: HCl(aq), HNO3(aq), HClO4(aq) and H2SO4(aq)* *for the first proton only

9. Acids & Bases Most acids are weak acids in which only a small percentage of the molecules dissociate to protonate water. HA(aq) + H2O(aq) ↔ H3O+(aq) + A-(aq) HA is a generic monoprotic weak acid such as HF, HCN or CH3COOH.

10. Weak Acids Weak acids in water form an equilibrium with hydronium ion and the deprotonated anion of the acid. The reaction does not go to completion. H-A + H2O ↔ H3O++ A-

11. Acids The equilibrium of acids in water can be viewed as a competition between the forward reaction and the reverse reaction. H-A + H2O ↔ H3O++ A-

12. Acids The forward reaction involves HA, a generic acid, and water. Water acts as a base by accepting a proton from the acid. H-A + H2O ↔ H3O++ A-

13. Acids The forward reaction involves HA, a generic acid, and water. Water acts as a base by accepting a proton from the acid. Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

14. Acids The reverse reaction involves A- accepting a proton and acting as a base. H3O+ donates a proton, and is an acid. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

15. Conjugate Acids & Bases The deprotonated acid is a base, and the protonated base is an acid. These acids and bases are called conjugate acids and bases. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

16. Acids For weak acids, the equilibrium lies to the left, indicating that A- is a stronger base than water. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor

17. Acids Acid HA and base A- are related, and differ only by the addition or removal of H+. Acid H3O+ Acid HA Base H2O Base A- remove H+ add H+

18. Acids HA and A- are called conjugate acid-base pairs. A- is the conjugate base of the acid HA. Acid H3O+ Acid HA Base H2O Base A- remove H+ add H+

19. Acids Likewise, H2O and H3O+ are related, and differ only by the addition or removal of H+. Acid H3O+ Acid HA Base H2O Base A- add H+ remove H+

20. Acids H2O and H3O+ are conjugate acid-base pairs. H2O is the conjugate base of H3O+. Acid H3O+ Acid HA Base H2O Base A- add H+ remove H+

21. Strong and Weak Acids HX is a weak acid and forms only a small amount of H3O+ and X-

22. Strong and Weak Acids HY is a strong acid and dissociates completely to form H3O+ and Y-

23. Strong and Weak Acids

24. Acid Strength The conjugate bases of strong acids have no tendency to pick up a proton. There is no reverse reaction. The conjugate bases of infinitely strong acids are infinitely weak.

25. Conjugate Acid-Base Strength The weaker the acid, the stronger its conjugate base. As the acid gets weaker, the reverse reaction with water becomes more significant.

26. Conjugate Acid-Base Strength

27. Ka Values Acid strength is determined by measuring the equilibrium constant for the following reaction: HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA]

28. Ka Values HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA] Water is left out of the equilibrium constant expression because it is a pure liquid (with constant concentration). The “a” subscript stands for acid.

29. Ka Values

30. Sulfuric Acid Sulfuric acid, H2SO4, can lose two protons. It is called a diprotic acid. Sulfuric acid is also one of the common strong acids. This applies to loss of the first proton only. H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq) This reaction goes 100% to the right.

31. Sulfuric Acid Sulfuric acid, H2SO4, can lose two protons. It is called a diprotic acid. Sulfuric acid is also one of the common strong acids. This applies to loss of the first proton only. H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq) HSO4- can react with water to lose an additional proton.

32. Sulfuric Acid HSO4-(aq) + H2O(l) ↔ H3O+(aq) + SO42-(aq) The double arrows indicate an equilibrium is established because HSO4- is a weak acid. Ka for HSO4- = [H3O+][SO42-] = 1.2 x 10-2 [HSO4-]

33. Sulfuric Acid HSO4-(aq) + H2O(l) ↔ H3O+(aq) + SO42-(aq) Ka for HSO4- = [H3O+][SO42-] = 1.2 x 10-2 [HSO4-] The value of Ka indicates that hydrogen sulfate ion is a relatively strong weak acid. It is stronger than most weak acids, but weaker than the strong acids HCl, HNO3, H2SO4 or HClO4.

34. Ka Values

35. Bases Strong bases are very effective at accepting protons. The most common strong bases are the soluble group IA and IIA metal hydroxides. In general, the metal ion is non-reactive, and serves as a spectator ion. Another strong base is oxide ion, O2-. Oxide reacts with water to become fully protonated. O2-(aq) + H2O(l)  2 OH-(aq)

36. Bases O2-(aq) + H2O(l)  2 OH-(aq) In this reaction, water is donating a proton, and hence acting as an acid. In previous reactions, water accepted a proton and served as a base. Substances that can behave as either an acid or a base are called amphoteric.

37. Amphoteric Nature of Water Depending upon its environment, water may donate a proton, acting as an acid, or accept a proton, and act as a base. This behavior is characteristic of amphoteric substances. Pure water molecules can react with each other, to a very small extent, to form hydronium and hydroxide ions.

38. Autoionization of Water This process is called autoionization or selfionization. One water molecule donates a proton to another. The result is the formation of equal amounts of hydronium and hydroxide ions. H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)

39. Autoionization of Water H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq) Kw = [H3O+][OH-] = 1.0 x 10-14 at 25oC In pure water, the concentration of hydronium ion equals the concentration of hydroxide. Both ions have a concentration of 1.0 x 10-7M.

40. Autoionization of Water

41. [H3O+] and [OH-] in Aqueous Solution The product of the hydroxide and hydronium concentration in any aqueous solution must equal Kw. As a result, when a solution is acidic, the hydronium concentration increases, and the hydroxide concentration decreases.

42. [H3O+] and [OH-] in Aqueous Solution Likewise, in basic solutions, the hydroxide ion concentration is greater than the hydronium ion concentration. There is always some hydroxide ion and some hydronium ion present in any aqueous solution.

43. [H3O+] and [OH-] in Aqueous Solution

44. The pH Scale A scale of acidity, the pH scale, is used to indicate the degree of acidity of aqueous solutions. pH = -log[H3O+] or –log[H+] The scale generally runs from 0-14, though negative pH values are possible. A neutral solution will have a pH = 7.00

45. The pH Scale A one unit change in pH is a ten-fold change in the concentration of hydronium ion. Acidic solutions have pH values less than 7.00, and basic solutions have pH values greater than 7.00

46. pH Values Most foods have pH values in the acidic range.

47. Problem: Calculation of pH • Calculate the pH of 0.10M HCl. - Is it an acid or a base? - Is it strong or weak? - Write the appropriate chemical reaction(s). HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) or HCl(aq)  H+(aq) + Cl-(aq)

48. Problem: Calculation of pH • Calculate the pH of 0.10M HCl. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) [H3O+] = 0.10M pH = - log [H3O+] = - log( 0.10) = 1.00 For every significant digit in the concentration, there is a place after the decimal in the pH.

49. Problem: Calculation of pH • Calculate the pH of 0.10M HCl. pH = - log [H3O+] = - log( 0.10) = 1.00 - Does your answer make sense? The pH of a fairly concentrated strong acid should be much less than 7. Yes, the answer makes sense.

50. Question: pH • Can you have a negative value for pH? Under what circumstances? If the hydronium concentration is > 1.0M, the pH will be negative. pH will be zero if the hydronium concentration equals 1.0M.