Acids & Bases CHM 1046 Bushra Javed Valencia College

1. Acids & Bases CHM 1046 BushraJaved Valencia College

2. Contents 1. Arrhenius Concept of Acids and Base 2. Brønsted–Lowry Concept of Acids and Bases 3. Lewis Concept of Acids and Bases 4. Relative Strengths of Acids and Bases 5. Molecular Structure and Acid Strength 6. Self-Ionization of Water 7. Solutions of a Strong Acid or Base 8. The pH of a Solution

3. Acids and Bases We will examine three ways to explain acid–base behavior: Arrhenius Concept: in terms of H+(aq) OH−(aq) Bronsted–Lowry Concept: in terms of donor & acceptor of a proton (H+) Lewis Concept: in terms of donor & acceptor of an electron pair

4. Arrhenius Concept of Acids and Bases An Arrhenius acid is a substance that, when dissolved in water, increases the concentration of hydronium ion, H3O+(aq). An Arrhenius base is a substance that, when dissolved in water, increases the concentration of hydroxide ion, OH-(aq).

5. Arrhenius Concept of Acids and Bases Example 1 Which of the following species is not capable of acting as an Arrhenius acid? a) H2SO3 b) SO32– c) H2O d) H3O+

6. Arrhenius Concept of Acids and Bases Example 2 Which of the following statements does not describe a characteristic of an Arrhenius acid? a) An Arrhenius acid is an electrolyte. b) An Arrhenius acid turns red litmus blue. c) An Arrhenius acid tastes sour. • An Arrhenius acid reacts with an Arrhenius base to produce a salt.

7. Arrhenius Concept of Acids and Bases Example 3 Which of the following species is not capable of acting as an Arrhenius base? a) NaOH b) KOH c) Ba(OH)2 d) NH3

8. Limitations of Arrhenius concept The Arrhenius concept limits bases to compounds that contain a hydroxide ion. The Bronsted–Lowry concept expands the compounds that can be considered acids and bases.

9. Bronsted–Lowry Concept of Acids and Bases An acid–base reaction is considered a proton (H+) transfer reaction.

10. Bronsted–Lowry Concept of Acids and Bases A Bronsted–Lowry acid is the species donating a proton in a proton-transfer reaction; it is a proton donor. A Bronsted–Lowry base is the species accepting a proton in a proton-transfer reaction; it is a proton acceptor.

11. Conjugate Acid–Base pair Substances in the acid–base reaction that differ by the gain or loss of a proton, H+, are called a conjugate acid–base pair. The acid is called the conjugate acid; the base is called a conjugate base. Conjugate acid Conjugate base Acid Base

12. Conjugate Acid–Base pair What is the conjugate acid of H2O? What is the conjugate base of H2O?

13. Conjugate Acid–Base pair Example 4 Identify the Bronsted–Lowry acids in the following equilibrium? HCOO–(aq) + H2O(l) ↔ HCOOH(aq) + OH–(aq) a) HCOO– and HCOOH b) HCOO– and OH– c) H2O and HCOOH d) H2O, HCOOH, and OH–

14. Conjugate Acid–Base pair Example 5 What is the conjugate base of H2PO4–(aq)? • H3O+ b) H3PO4 c) HPO42– d) PO43–

15. Conjugate Acid–Base pair Example 6 Label each species as an acid or base. Identify the conjugate acid-base pairs. a. HCO3−(aq) + OH−(aq) ↔CO32−(aq) + H2O(l) b. HCO3−(aq) + HF(aq) ↔H2CO3(aq) + F−(aq)

16. Amphoteric species Species that can act as both an acid and a base are called amphiproticor amphoteric species. Identify any amphiprotic species in the previous problem. HCO3−was a base in the first reaction and an acid in the second reaction. It is amphiprotic.

17. Lewis Concept of Acids and Bases A Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species. A Lewis base is a species that can form a covalent bond by donating an electron pair to another species.

18. Lewis Concept of Acids and Bases Example 7 Which of the following species cannot act as a Lewis base? a) S2– b) SH– c) Al3+ d) H2O

19. Lewis Concept of Acids and Bases Example 8 Which of the following species cannot act as a Lewis acid? a) K+ b) Mg2+ c) Al3+ d)H–

20. Relative Strengths of Acids and Bases

21. Relative Strengths of Acids and Bases The stronger an acid, the weaker its conjugate base. The weaker an acid, the stronger its conjugate base.

22. Relative Strengths of Acids and Bases Example 9 Which of the following acids has the strongest conjugate base? a) HClO4 b) HClO3 c) HClO2 d) HClO

23. Relative Strengths of Acids and Bases Example 10 Which of the following acids has the weakest conjugate base? a) HF b) HI c) CH3COOH d) HNO2

24. Molecular Structure and Acid Strength The strength of an acid depends on how easily the proton, H+, is lost or removed. • Therefore, more polarized the bond between H and the atom to which it is bonded, the more easily the H+ is lost or donated.

25. Molecular Structure and Acid Strength For a binary acid, as the size of X in HX increases, going down a group, acid strength increases For a binary acid, going across a period, as the electronegativity increases, acid strength increases

26. Molecular Structure and Acid Strength Example 11 Which of the following statements is incorrect? • One reason why HCl is a stronger acid than HF is that Cl has a larger atomic radius than F. • One reason why HCl is a stronger acid than HF is that the H–Cl bond is weaker than the H–F bond. • One reason why HCl is a stronger acid than HF is that Clis more electronegative than F. • The acids HBr and HI both appear equally strong in water.

27. Molecular Structure and Acid Strength H2O and H2S These are binary acids from the same group, so we compare the size of O and S. Because S is larger, H2S is the stronger acid. HCl and H2S These are binary acids from the same period, but different groups, so we compare the electronegativity of Cl and S. Because Cl is more electronegative, HCl is the stronger acid.

28. Acid Strength of Oxoacids For a series of oxoacids having the same central atom, acid strength increases as number of oxygensincreases. Example 12 Which of the following acids has the strongest conjugate base? a) HClO4 b) HClO3 c) HClO2 d) HClO

29. Acid Strength of Oxoacids For a series of oxoacids differing only in the central atom , the acid strength increases with the electronegativity of central atom

30. Acid Strength of Oxoacids Example 13 The acid strength decreases in the series HBr > HSO4– > CH3COOH > HCN > HCO3–. Which of the following is the strongest base? a) CO32– b) CN– c) CH3COO– d) SO42–

31. Acid strength of polyproticacids The acid strength of a polyprotic acid and its anions decreases with increasing negative charge. H2CO3 is a stronger acid than HCO3−. H2SO4 is a stronger acid than HSO4−. H3PO4 is a stronger acid than H2PO4−. H2PO4- is a stronger acid than HPO42

32. Direction of Reaction A reaction will always go in the direction: • from stronger acid to weaker acid • and from stronger base to weaker base.

33. Direction of Reaction Let’s see which species are favored at the completion of the following reaction: HCN(aq) + HSO3−(aq) ↔ CN−(aq) + H2SO3(aq) • We first identify the acid on each side of the reaction: HCN and H2SO3. • Next, we compare their acid strength: H2SO3 is stronger. • This reaction will go from right to left (), and the reactants are favored.

34. Direction of Reaction Example 14 Which of the following reactions is not product-favored? a) HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq) b) HClO4(aq) + H2O(l) → H3O+(aq) + ClO4–(aq) c) NH3(aq) + H2O(l) → NH4+(aq) + OH–(aq) d) NaOH(aq) → Na+(aq) + OH–(aq)

35. Direction of Reaction Example15 The reaction of which acid with water is product-favored? a) nitric acid b) chloric acid c) nitrous acid d) chlorous acid

36. Self-Ionization of Water H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq) Base Acid Conjugate acid Conjugate base

37. Self-Ionization of Water H2O(l) + H2O(l)↔H3O+(aq) + OH−(aq) We call the equilibrium constant the ion-product constant, Kw. Kw = [H3O+][OH−] At 25°C, Kw = 1.0 × 10−14 As temperature increases, the value of Kw increases.

38. Solutions of a Strong Acid or Strong Base The concentration of hydronium or hydroxide in a solution of strong acid or base is related to the stoichiometry of the acid or base.

39. Solutions of a Strong Acid or Strong Base Example16 Calculate the hydronium ion concentration at 25°C in 0.10 MHCl When HCl ionizes, it gives H+ and Cl−. So [H+] = [Cl−] = [HCl] = 0.10 M.

40. Solutions of a Strong Acid or Strong Base Example 17 Calculate the hydroxide ion concentration at 25°C in 1.4 × 10−4MMg(OH)2 • When Mg(OH)2 ionizes, it gives Mg2+ and 2 OH−. • So [OH−] = 2[Mg2+] = 2[Mg(OH)2] = 2.8 × 10−4 M.

41. Aqueous Solutions Solutions can be characterized as Acidic: [H3O+] > 1.0 × 10−7M Neutral: [H3O+] = 1.0 × 10−7M Basic: [H3O+] < 1.0 × 10−7M

42. The pH Concept • pH is a measure of the acidity of a solution. • A neutral solution has a pH of 7, an acidic solution has a pH less than 7, and a basic solution has a pH greater than 7. • The pH scale uses powers of 10 to express the hydrogen ion concentration.

43. The pH Concept

44. The pH of a Solution pH = –log[H3O+] For a log, only the decimal part of the number has significant figures. [H3O+] = 10−pH

45. The pH of a Solution pOH= –log[OH−] [OH−] = 10−pOH pH + pOH = 14.00 (at 25°C)

46. The pH of a Solution Example18 The pH of natural rain in 5.60. What is its hydronium ion concentration pH = 5.60 [H3O+] = 10−pH = 10−5.60 [H3O+] = 2.5 × 10-6 M Because the pH has two decimal places, the concentration can have only two significant fig

47. The pH of a Solution Example19 The concentration of H3O+ in a solution is 8 × 10-4M at 25ºC. What is its hydroxide-ion concentration? a) 8 × 10-4M b) 1 × 10-10M c) 2 × 10-10M d)1 × 10-11M

48. The pH of a Solution Example 20 A solution has a hydroxide-ion concentration of 7.48 × 10–5M. What is its hydronium-ion concentration? a) 1.00 × 10–7M b) 1.34 × 10–10M c) 7.48 × 10–5M d) 7.48 × 10–19M