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Chapter 17: Chemistry of Acids & Bases

Chapter 17: Chemistry of Acids & Bases. “The end is near”. 17.0 Objectives. 1. To recognize the properties of strong and weak acids and strong and weak bases. 2. To become familiar with Arrhenius and Bronsted-Lowry theories of acids and bases.

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Chapter 17: Chemistry of Acids & Bases

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  1. Chapter 17: Chemistry of Acids & Bases “The end is near”

  2. 17.0 Objectives • 1. To recognize the properties of strong and weak acids and strong and weak bases. • 2. To become familiar with Arrhenius and Bronsted-Lowry theories of acids and bases. • 3. Be able to calculate and understand the following quantities: [H+], [OH-], pH, pOH, Kw, Ka, and Kb. • 4. Calculate the equilibrium concentrations of weak acids and bases. • 5. Classify aqueous salt solutions as acidic, basic, or neutral. • 6. Determine the concentration of unknown acid or base, or the Ka of a weak acid from a titration procedure.

  3. Homework • HW #1 - 11, 13, 15, 17, 19, 21 • Conjugate Acids/Bases; pH • HW #2 - 25, 27, 29, 31, 33, 37, 89 • Equilibrium Constants and pKa • HW#3 - 39, 41, 43, 45, 47, 49 • pKb, Types of Reactions • HW#4 - 51, 53, 55, 93 • pH to calculate K values • HW#5 - 57, 59, 63, 65, 97, 101 • Equilibrium problems

  4. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • A. Definitions • 1. Arrhenius definition of acid • Substance that contains H+ • Ex. HCl; H2SO4 • 2. Arrhenius definition of base • Substance that contains OH- • Ex. NaOH, Ca(OH)2

  5. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • B. Properties of acids • Ionize when put into water • React with active metals (Group I, II) to produce Hydrogen gas • Neutralized with bases • Have a sour taste

  6. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • C. Properties of bases • Ionize when put into water • Neutralized with acids • Have a bitter taste • Feel Slippery

  7. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • D. Strong acids - definition and examples • Strong acids dissociate completely in aqueous solution: • Strong Acids • HCl, HBr, HI, HNO3, HClO4, H2SO4 • “BrIClSONOClO 4 3 4” • EX. HCl

  8. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • Weak acids - definition • Weak acids are partly dissociated in soln. • Treated as an equilibrium • Ex. CH3COOH • Species • H+ • H3O+ • [insert H-containing cation here]+

  9. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • F. Example 17.1 Write equations for the aqueous dissociation of the following weak acids: HPO4-2, Fe(H2O)63+, HCN

  10. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • G. Strong bases - definition and examples • Strong bases completely dissociate in aqueous solution: • Group I & MOST Group II Hydroxides • LiOH, NaOH, KOH, RbOH, CsOH • Ca(OH)2, Sr(OH)2, Ba(OH)2, • Ex. NaOH

  11. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • H. Weak bases - definition and species • Two types of substances act like weak bases in aqueous solution: • Nitrogen-containing compounds • Ex. NH3 • Anions of acids • Ex. HCO3-

  12. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • I. Example 17.2 Write equations for the water dissociation of the following weak bases: NO21-, CO32-, C2H5NH2

  13. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • A. Arrhenius definitions – problems • Requires that the solution be aqueous in order to work • Constrains acids and bases to contain H+ and OH-

  14. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • B. Bronsted-Lowry acid - definition and example • Acid = Proton Donor • “For every acid, there must be a base” • Ex. HNO3

  15. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • C. Bronsted-Lowry base - definition and example • Base= Proton acceptor

  16. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • D. Conjugate acids and bases • Conjugate acid: Bronsted Base with an extra proton (H+) attached • Every base has a conjugate acid • Conjugate base: Bronsted acid minus a proton (H+) • Every acid has a conjugate base

  17. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • E. Amphiprotic substances • Substances that can ionize as either an acid or a base depending on the properties of the other species in soln. • Ex. H2O • can act as a conjugate acid or a conjugate base

  18. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • F. Example 17.3 Identify the Bronsted Lowry acids and bases and their conjugates in the following equations: • NH4+(aq) + OH1-(aq)  NH3(aq) + HOH(l) • N2H4(aq) + HOH(l)  N2H51+(aq) + OH1-(aq)

  19. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • G. Example 17.4 Write equations to show that HSO41- can act like an amphiprotic substance.

  20. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • H. Polyprotic acids • Acids that give off more than 1 H+ when put into water • 1st Proton usually given off rapidly • Subsequent protons are given off with increasing difficulty (stronger bases at each step)

  21. 17.3WATER AND THE pH SCALE • A. Water - self-ionization, and the water dissociation constant, Kw H2O + H2O --- H3O+ + OH- Acid(1) Base(2) Acid(2) Base(1) • For any sample of water molecules: • 2 H2O (l) <-----> H3O+ (aq) + OH- (aq) • Kw = = 1.0 x 10-14

  22. 17.3WATER AND THE pH SCALE • B. Relationships for all aqueous solutions • In a neutral soln: • [H3O+] = [OH-] = 1.0 x 10-7 M • In an acidic soln: [H3O+] > [OH-] • [H3O+] > 1.0 x 10-7 M • [OH-] < 1.0 x 10-7 M • In a basic soln: [H3O+] < [OH-] • [H3O+] < 1.0 x 10-7 M • [OH-] > 1.0 x 10-7 M

  23. 17.3WATER AND THE pH SCALE • C. The pH scale • Ranges from 0 to 14 • 0~7 = ACID • 7~14 = BASE • 7 = Neutral

  24. 17.3WATER AND THE pH SCALE • D. Definition of pH and pOH • pH = -log [H+] • [H3O+] = 10-pH • Eqns. for pOH? • pH + pOH = 14

  25. 17.3WATER AND THE pH SCALE • E. Example 17.5 Calculate the [H+], [OH1-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4

  26. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • A. Weak acids, Ka • Equilibrium of an acid with conjugate base • **REMEMBER** H2O (l) is not included! HA + H2O <-----> H3O+ + A- Ka =

  27. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • B. Value • Less than 1 because it is a weak acid (more reactant compared to product) • As Ka increases the strength of acid increases (directly related)

  28. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • C. pKa and value • pKa = - log Ka • As pKa decreases the strength of the acid increases (inversely related)

  29. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • D. Example 17.6 In a solution prepared by dissolving 0.100 mole of lactic acid per liter, [H+] = 3.7 x 10-3M. Calculate the Ka for lactic acid.

  30. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • E. Example 17.7 Determine the [H+], pH and % dissociation of carbonic acid, H2CO3 (Ka = 4.2 x 10-7) in a 0.100 M solution.

  31. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • F. Approximation • Small Ka values = few protons given off = very little dissociation • If Ka is smaller than 10-3, “x” is negligible in an IRE where it is “-x” • For quadratic approximations only!

  32. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • Weak Bases, Kb • Equilibrium of a base with conjugate acid • **REMEMBER** H2O (l) is not included! B- + H2O <-----> OH- + HB Kb =

  33. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • pKb and Kb values • pKb = - log Kb • As pKb decreases the strength of the base increases (inversely related) • Large pKb = small pKa • Ka x Kb = Kw = 1.0 x 10-14 • pKa + pKb = pKw = 14.0 • Think of Ka and Kb as opposites

  34. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • I. Example 17.8 Pyruvic acid, HPy, has a Ka of 3.3 x10-3 and Butyric acid, HBut, has a Ka of 2.0 x 10-5. Which is the stronger acid, HPy or HBut? Which is the stronger base, But1- or Py1-?

  35. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • J. Example 17.9 Determine the [H+], [OH1-], pH and pOH of a 0.100M solution of a weak base that has a Kb of 1.0 x 10-9.

  36. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • K. Acid –Base Properties of salt solutions • salt – definition • “A compound formed by replacing the hydrogen ions of an acid with metal ions”

  37. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • 2. salt dissociation in water – hydrolysis reaction • Many of the ionizations we have been writing are hydrolysis • Ex. Fe(H2O)6+3 weak acid

  38. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • 3. interaction of ions with water molecules • Anions that are conjugate bases of strong acids are extremely weak bases • Alkali/Alkaline earth metals have no effect

  39. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES

  40. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • 4. Example 17.10 For the following salts, determine whether the aqueous solution would be acidic or basic or neutral. Then write the net ionic equation for the interaction of each active ion with water. KClO4 Na2HPO4 NH4I

  41. 17.6 TYPES OF ACID-BASE REACTIONS • 1. Strong acid and strong base – neutralization • 2. weak acid and strong base • 3. weak base and strong acid • 4. weak base and weak acid

  42. Strong acid and strong base - neutralization

  43. Weak acid and strong base

  44. Weak base and strong acid

  45. Weak base and weak acid

  46. 17.6 TYPES OF ACID-BASE REACTIONS • 5. Example 17.11 Write net ionic equations for the interactions between the following pairs of acids and bases: a. HNO2 and NaOH b. (CH3)2NH and HCl c. HBr and KOH.

  47. 17.6 TYPES OF ACID-BASE REACTIONS • 5. Acid Base Titrations • Complete dissociations • No need to worry about Ka or Kb • Strong Acid + Strong Base --> H2O + Salt Equation • (#H+) Ma Va = Mb Vb (#OH-) • Stoich will also work 

  48. 17.6 TYPES OF ACID-BASE REACTIONS • B. Equivalence point and general graph • Point in a titration at which one reactant has been exactly consumed by adding the other reactant • GRAPH:

  49. 17.6 TYPES OF ACID-BASE REACTIONS • 6. Example 17.12 A chemist synthesizes a sample of nicotinic acid (HC6H4NO2 molar mass 123g/mole). To determine the purity of the sample, a 0.450g portion of the acid is titrated with 0.100M NaOH. A total volume of 35.8mL of NaOH is required to neutralize the nicotinic acid according to the following: • HC6H4NO2(aq) + OH1-(aq)  HOH(l) + C6H4NO21-(aq) • Determine the percentage of nicotinic acid in the original sample.

  50. 17.6 TYPES OF ACID-BASE REACTIONS • 7. Example 17.13 If 50.00mL of 1.000M HCl is titrated with 1.000M NaOH, find the pH of the solution after the following volumes of NaOH have been added: 25.00mL, 50.00mL, and 50.01mL.

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