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. Lewis Model • More general definition than either Arrhenius or Bronsted-Lowry • Includes substances that are not classified as acids or bases under the other definitions • Acid - a substance that can accept a pair of electrons to form a covalent bond • Base - a substance that can donate a pair of electrons to form a covalent bond

11. 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-

12. 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

15. Hydronium Ions - H3O+ ion • In water - H+ ions are strongly attracted to the electrons surrounding water molecules • When one compound in a reaction acts as an acid (donate an H+ ion) the other acts as a base (accepts an H+ ion)

16. 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

17. Strengths of Acids and Bases • Strong acid/base – acids/bases that dissociate completely in water  • Strong Acids - HCl, H2SO4, HBr , HNO3 HI , HClO4 • Strong Bases – LiOH, NaOH, Ca(OH)­2 KOH, Sr(OH)2 , RbOH, Ba(OH)2 , CsOH

18. Strengths of Acids and Bases • Weak acid/base – acids/bases that dissociate only partly in water • There is an inverse relationship between the strengths of conjugate acid-base pairs • The stronger the acid, the weaker the conjugate base and vise versa • The stronger the base, the weaker the conjugate acid and vise versa

19. 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

20. 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

21. Ka Example • Complete the reaction and write the equilibrium constant expression for the following reactions • HCl (aq) + H2O (l) • HCO3- (aq) + H2O (l)

22. Ka Example • A monoprotic weak acid has a concentration of 0.092M. At equilibrium, the concentration of hydronium is 0.0024M. What is the Ka for this acid?

23. 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.

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. Kb Example • Complete the reaction and write the equilibrium constant expression for the following reactions • CH3NH2 (aq) + H2O (l) • NH4+ (aq) + H2O (l)

26. Kb Example • 0.23mol of a weak base is mixed with 1.5L of water. At equilibrium, the concentration of OH- is 0.0040M. What is the Kb for the base?

27. Salt Hydrolysis Reactions • Reactions of ions from salts to form H3O+or OH- ions • Can predict whether a salt hydrolysis rxn produces an acidic sol’n or basic sol’n • Consider the acid and base the salt is formed from – see next slide

28. Salt Hydrolysis Reactions • 4 possibilities • strong acid + strong base = neutral salt • strong acid + weak base = acidic salt • weak acid + strong base = basic salt • weak acid + weak base = acidic, basic, or neutral salt, depending on the relative strengths of the acids and bases which the salt is formed

29. Identifying Acids • Acids • Acidic hydrogen – hydrogens that can be donated • Not every hydrogen is acidic - HC2H3O2 • Acidic hydrogens - have a slight positive charge while it is still a part of a molecule • the acidic hydrogen is on the positive end of a polar covalent bond

30. 3 Categories of Acids • binary acids – contains hydrogen and 1 other element • oxy acids – contains hydrogen, oxygen, and 1 other element (polyatomic ions) • carboxylic acids – organic acids; acids that contain carbon atoms

31. Identifying Bases • Bases • Bronsted-Lowry base always contains an unshared pair of electrons • 2 categories of bases • anions – includes monatomic and polyatomic anions • amines – contains nitrogen atoms that has an unshared pair of electrons; ammonia derivative

32. Ex. What is the concentration of hydroxide ions in blood, if the hydronium ion concentration is 4.5 x 10-8? Is blood acidic, basic, or neutral?

33. 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

34. 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

35. 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

36. Measuring pH • 2 ways to measure pH • indicators • pH meters • both detect the presence of H3O+ ions • indicators change color based on the H3O+ ions • common indicators: litmus paper, thymol blue, methyl orange, methyl red, bromthymol blue, phenolphthalein

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

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

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

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

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

42. Buffers • A mixture that is able to release or absorb H+ ion, keeping a solutions pH constant • Can control the pH within very narrow limits • Most common buffers are mixtures of weak acids and their conjugate bases

43. Buffer Capacity • The amount of acids or base that a buffer can neutralize • All buffers have a limited capacity to neutralize H3O+ ions or OH- ions • If you add H3O+ ions or OH‑ ions beyond the buffer capacity, the ions will remain in the solution, changing the pH • The greater concentration of buffer in the solution, the greater the buffer capacity

44. 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

45. 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

46. 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

47. 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

49. 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

50. 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?