Acids and Bases

1. Acids and Bases

2. Properties of Acids • Sour taste • React w/ metals to form H2 • Most contain hydrogen • Are electrolytes • Change color in the presence of indicators (turns litmus red) • Has a pH lower than 7

3. Two Types of Acids • Strong acids • Any acid that dissociates completely in aqueous sol’n • Weak acids • Any acid that partially dissociates in aqueous sol’n

4. Properties of Bases • Bitter taste • Slippery feel • Are electrolytes • Change color in the presence of indicators (turns litmus blue) • Has a pH higher than 7

5. Types of Bases • Strong Base • Any base that dissociates completely in aqueous sol’n • Weak Base • Any base that partially dissociates in aqueous sol’n

6. Neutralization • Neutralization rxn: a rxn of an acid and a base in aqueous sol’n to produce a salt and water • Salt: compound formed from the positive ion of a base and a negative ion of an acid • Properties of the acid and base cancel each other

7. Arrhenius Model of Acids and Bases • Proposed the model in 1887 • Acid: any compound that produces H+ ions in aqueous (water) sol’n • Base: any compound that produces OH- (hydroxide) ion in aqueous sol’n • Offers an explanation of why acids and bases neutralize each other (H+ + OH- = H2O)

8. Problems with Model • Restricts acids and bases to water sol’ns (similar reactions occur in the gas phase) • Does not include certain compounds that have characteristics of bases (e.g., ammonia)

9. Brønsted-Lowry Model of Acids and Bases • Brønsted acid: a hydrogen ion donor (H+, or proton) • Brønsted base: a hydrogen ion acceptor • Defines acids and bases independently of how they behave in water • Amphiprotic: having the property of behaving as an acid and a base • Also called amphoteric, e.g., water

10. Conjugate Acid-Base Pairs • The rxn between Brønsted-Lowry acids and bases can proceed in the reverse direction (reversible reactions) HX (aq) + H2O (l) H3O+ (aq) + X- (aq) • The water molecule becomes a hydronium ion (H3O+), and is an acid because it has an extra H+ to donate • The acid HX, after donating the H+, becomes a base X-

11. Conjugate Acids and Bases HX (aq) + H2O (l)  H3O+ (aq) + X- (aq) Base Conjugate Acid Acid Conjugate Base Forward reaction: Acid and base Reverse reaction: Conjugate acid and conjugate base

12. Conjugate Acid: species produced when a base accepts a hydrogen ion from an acid • Conjugate Base: species produced when an acid donates a hydrogen ion to a base • Conjugate Acid-Base Pair: two substances related to each other by the donating and accepting of a single hydrogen ion

15. Types of Acids • Monoprotic acids: acids that contain only 1 hydrogen; e.g., HCl • Diprotic acids: acids that contain 2 hydrogens; e.g. H2CO3 • Triprotic acids: acids that contain 3 hydrogens; e.g. H3PO4

16. More Types of Acids • Binary acids: acids that contain only 2 elements; e.g. HF • Polyatomic acids: acids that contain more than 2 elements; e.g. H2SO4 • These acids contain polyatomic ions • Also called ternary or oxy- acids

17. Naming Binary Acids • Start with the prefix hydro- • Put it in front of the root word of the anion (- charged ion) • Add –ic to the end • Examples • Hydrobromic (HBr) • Hydrofluoric (HF) • Hydroiodic (HI) • Hydrochloric (HCl)

18. Naming Polyatomic Acids • Start with the root word of the name of the polyatomic ion • Add –ous if name ends in –ite • Add -ic if name ends in –ate • Examples: • Chlorous (from chlorite, ClO2-) • Nitric (from nitrate, NO3-) • Sulfurous (from sulfite, SO3-2)

19. pH and [H3O+] • pH: number that is derived from the concentration of hydronium ions ([H3O+]) in sol’n • pH = -log [H3O+] • As pH increases, [H3O+] decreases • Scale ranges from 0 – 14 • pH = 7 is neutral • pH < 7 is acidic • pH > 7 is basic

20. p[OH] • pOH = - log [OH-] • pH + pOH = 14.00 • Calculating ion concentrations from pH • [H+] = antilog (-pH) • [OH-] = antilog (-pOH)

21. Dissociation Constants • Acid dissociation constant: (Ka): the equilibrium constant for the rxn of an aqueous weak acid and water • Base dissociation constant: (Kb): the equilibrium constant for the rxn of an aqueous weak base w/ water • Both are derived from the ratio of the concentration of the products and reactants at equilibrium

22. Acid Dissociation Constant Ka = [H3O+] [A-] [HA] • Ka is a measure of the strength of an acid • Ka values for weak acids are always less than one • Used mostly w/ weak acids because the Ka values for strong acids approach infinity

23. Examples • HMnO4 (aq) + H2O (l)  • H2S (aq) + H2O (l) 

24. Base Dissociation Constant Kb = [HB+] [OH-] [B] • Kb is a measure of the strength of a base • Kb values for weak bases are always less than 1 • Kb values for strong bases approach infinity

25. Examples • H2NOH (aq) + H2O (l)  • NH3(aq) + H2O (l)

26. Water • Water can dissociate into its component ions, H+ and OH- • 2H2O (l)  H3O+ (aq) + OH- (aq) • One water molecule acts as a weak acid, and the other acts as a weak base • The ions are present in such small amounts they can’t be detected by a conductivity apparatus • In pure water, [H3O+] =1.0 x 10 –7 M and [OH-] = 1.0 x 10-7 M

27. Dissociation Constant for Water • It is defined as Kw: the ion product constant for water • Kw = [H3O+] [OH-] • Kw = (1.0 x 10-7)(1.0 x 10-7) • Kw = 1.0 x 10-14 • The value of Kw can always be used to find the concentration of either H3O+ or OH- given the concentration of the other

28. Examples What is the pH of a 0.001 M sol’n of HCl, a strong acid?

29. Examples What is the pH of a sol’n if [H3O+] = 3.4 x 10-5 M?

30. Examples The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H3O+]?

31. Examples The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H3O+]?

32. Calculating Ka • In these problems, remember that the concentration of the [H3O+] ions will equal the concentration of the conjugate base ions. • This is because for every molecule of weak acid that dissociates, there will be an equal number of H3O+ ions and base ions

33. Example Assume that enough lactic acid is dissolved in sour milk to give a solution concentration of 0.100 M lactic acid. A pH meter shows that the pH of the sour milk is 2.43. Calculate Ka for the lactic acid equilibrium system.

34. Titrations • An analytical procedure used to determine the concentration of a sample by reacting it with a standard sol’n • In a titration, an indicator is used to determine the end point • Standard sol’n: a sol’n of precisely known concentration • Indicator: any substance in sol’n that changes color as it reacts with either an acid or a base

35. Titrations • Each indicator changes its color over a particular range of pH values (transition interval) • An unknown acid sol’n will be titrated with a standard sol’n that is a strong base • An unknown base sol’n will be titrated with a standard sol’n that is a strong acid

36. Titrations • Equivalence point: point at which the concentration of H3O+ ions is the same as the concentration of OH- ions; [H3O+ ] = [OH-] • Endpoint: the point at which the indicator changes color • Titration curve: graph that shows how pH changes in a titration

37. Titrations • The equivalence point is at the center of the steep, vertical region of the titration curve • At the equivalence point, pH increases greatly w/ only a few drops

39. Example Problem 1 What is the molarity of a CsOH solution if 20.0 mL of the solution is neutralized by 26.4 mL of 0.250M HBr solution? HBr + CsOH → H2O + CsBr

40. Example Problem 2 What is the molarity of a nitric acid solution if 43.33 mL 0.200M KOH solution is needed to neutralize 20.00 mL of unknown solution?

41. Example Problem 3 What is the concentration of a household ammonia cleaning solution (NH4OH) if 49.90 mL of 0.5900M H2SO4 is required to neutralize 25.00 mL solution?