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Chapter 17: Chemistry of Acids & Bases

Chapter 17: Chemistry of Acids & Bases

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Chapter 17: Chemistry of Acids & Bases

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  1. Chapter 17: Chemistry of Acids & Bases

  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 • Equilibrium Constants and pKa • HW#3 - 39, 41, 49 • pKb, Types of Reactions • HW#4 - 51, 53, 55 • pH to calculate K values • HW#5 - 57, 59, 97, 101 • Equilibrium problems • Mistake in Ch. 17 packet- too many Homework Problems Listed!!!! • YOU CAN SKIP CH 17 HW # 19, 89, 43, 45, 47, 93, 63, 65!! These should have been omitted!

  4. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • A. Definitions • 1. Arrhenius definition of acid • produce H+ in water solution • Ex. HCl; H2SO4 HCl H+ + Cl- • 2. Arrhenius definition of base • produce OH- in water solution • Ex. NaOH, Ca(OH)2 NaOH  Na+ + OH-

  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 • Neutralize bases to form water and salt • Have a sour taste • Turns blue litmus red

  6. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • C. Properties of bases • Ionize when put into water • Neutralize acids to form water and salt • Have a bitter taste • Feel Slippery • Turn red litmus blue

  7. Definition: Hydronium Ion In aqueous solution, H+ does NOT exist! Note: In problems, [H+] = [H3O+] H+ + H2O  H3O+ (hydronium ion) Dr. Mihelcic AP Chemistry 1

  8. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • D. REVIEW: Strong acids • Strong acids dissociate completely in aqueous solution: There are 6 Strong Acids • HCl, HBr, HI, HNO3, HClO4, H2SO4 • Examples: HCl  H+ + Cl- HNO3 H+ + NO3- Note single arrow!! No Equilibrium Established!

  9. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT Weak acids are much less than 100% ionized in water. (use double arrow; equilibrium established) • Weak acids • Weak acids are partly dissociated in soln. • Treated as an equilibrium • Ex. CH3COOH acetic acid • Species • H+ • H3O+ • [insert H-containing cation here]+

  10. 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

  11. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • G. Strong Bases: Review • Strong bases completely dissociate in aqueous solution: • Group I metal hydroxides LiOH, NaOH, KOH, RbOH, CsOH • MOST Group II Hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2, Examples: NaOH - Sr(OH)2

  12. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT Weak bases are less than 100% ionized in water.(use double arrow; equilibrium established) • H. Weak bases • partly dissociated in soln • Treated as an equilibrium • Two types of substances act like weak bases in aqueous solution: • Nitrogen-containing compounds (why?) • Ex. NH3 NH3(aq) + H2O(liq) NH4+(aq) + OH-(aq) • Anions of acids • Ex. F- , HCO3-

  13. Strong Acids and bases are strong electrolytes; weak acids and bases are weak electrolytes!

  14. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • I. Example 17.2 Write equations for the aqueous dissociation of the following weak bases: NO21- CO32- C2H5NH2

  15. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • A. Arrhenius definitions are Narrow Problems: • Only Includes aqueous systems • Limits acids and bases to contain H+ and OH-

  16. 17.2 BRONSTED CONCEPT OF ACIDS AND BASESA more general definition than Arrehenius theory… • B. Bronsted-Lowry Acid • ACIDS DONATE H+ IONS • Acid = Proton Donor (ADP: Acids Donate Protons) • Ex. HNO3 • C. Bronsted-Lowry Base • BASES ACCEPT H+ IONS • Base= Proton acceptor (BAP: Bases Accept Protons)

  17. The Brønsted definition means NH3 is a BASEin water, and water is itself an ACID Bases Accept Protons Acids Donate Protons

  18. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • D. Conjugate acids and bases • Conjugate acid: Bronsted Base w/ extra proton (H+) attached • Every base has a conjugate acid formed when H+ is added to the base. • Conjugate base: Bronsted acid minus a proton (H+) • Every acid has a conjugate baseformed when H+ is removed from the acid. • Conjugate Pairs: NH3 / NH4+ is a conjugate pair • related by the gain or loss of H+

  19. Conjugate Pairs Generalized equation: HB (aq) + A (aq) HA+(aq) + B-(aq)

  20. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES Some substances can function as both an ACID OR a BASE, depending on what they are reacted with. (can donate OR accept H+) • 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 H2O + H2O H3O+ + OH- H2O +B  OH-+ BH+ H2O + AH  H3O++ A-

  21. 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)

  22. 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) Acid Base Conj. Base Conj acid Bases Accept Protons Acids Donate Protons Acid Conj acid Base Acid Conj. Base Bases Accept Protons Acids Donate Protons

  23. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES HSO41-  • G. Example 17.4 Write equations to show that HSO41- can act like an amphiprotic substance. HSO41- 

  24. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • G. Example 17.4 Write equations to show that HSO41- can act like an amphiprotic substance. HSO41-  SO42- + H+ or HSO41- + H2O  SO42- + H3O+ HSO41- + H+  H2SO4 or HSO41- + H2O H2SO4 + OH-

  25. 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- why?) • Ex. H2SO4 HSO4- SO42- Another example: Oxalic acid (diprotic)

  26. Oxalic acid is diprotic, having two protons to donate Another example: phosphoric acid (Triprotic) Donates one H+ at a time: H3PO4 H+ + H2PO4- H2PO4-  H+ + HPO4-2 HPO4-2  H+ + PO4-3

  27. 17.3 WATER AND THE pH SCALE • A. Water -self-ionization, and the water dissociation constant, Kw Self-Ionization: (H2O ↔ H+ + OH-) 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 = [H3O+] [OH-] = 1.00 x 10-14 (at 25 oC)

  28. 17.3 WATER 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 2 H2O (l) ↔ H3O+ (aq) + OH- (aq) Remember: Kw = [H3O+] [OH-] = 1.00 x 10-14

  29. A common way to express acidity and basicity is with pH (the “power of hydrogen”) 17.3 WATER AND THE pH SCALE • C. The pH scale (Ch. 5 review) • Ranges from 0 to 14 • 0~7 = ACID • 7~14 = BASE • 7 = Neutral

  30. 17.3 WATER AND THE pH SCALE • D. Definition of pH and pOH • pH = - log [H+] • [H3O+] = 10-pH • pOH = - log [OH-] • [OH-] = 10-pOH • pH + pOH = 14 Remember, pure water (neutral) [H3O+] = [OH-] = 1.0 x 10-7 M

  31. pH 0: [H+] = 0.1 M pH 14: [H+] =0.00000000000001 M pOH 14: [OH-]=0.1 M pOH 0: [OH-]=0.00000000000001 M

  32. Simple log examples: • Definition of pH and pOH • pH = - log [H+] • [H+] = 10-pH • or: [H+] = antilog(-pH) • pOH = - log [OH-] • [OH-] = 10-pOH[OH-] = 10-pOH • or:[OH-] = antilog(-pOH) • pH + pOH = 14 [H3O+] + [OH-] = 1.0 x 10-14 M

  33. pH to [H+] Calculations pH of Coke is 3.12 (acidic) pH = - log [H3O+] log [H3O+] = - pH Take antilog and get [H3O+] = 10-pH [H3O+] = 10-3.12 = 7.6 x 10-4 M

  34. 17.3 WATER AND THE pH SCALE [H+]=5.0 x 10-3M pH = 7.4 • E. Example 17.5 Calculate the [H+], [OH-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4. pH = - log [H+] Kw= [H3O+] [OH-] =1 x10-14 pH + pOH = 14 10-pOH = [OH-]

  35. pH= - log [H+] pH + pOH = 14 10-pOH = [OH-] • E. Example 17.5 Calculate the [H+], [OH-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4.

  36. Note: Strong Acids and Bases Since strong Acids dissociate completely in aqueous solution: E.g. HCl  H+ + Cl- [H+] can be calculated from the molarity of the acid. 2.0 M 2.0 M 2.0 M Likewise, Strong Bases dissociate completely in aqueous solution: E.g. Sr(OH)2 Sr+2 + 2OH- So [OH-] can be calculated from the molarity of the base. 0.5 M 0.5 M 1.0 M

  37. Note: Weak Acids and Weak Bases- exixist in solution in equilibrium with conjugate Since srong Acids/bases do ionize less than 100% , we Must use ICE tables for all calculations involving weak acids!

  38. 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=

  39. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • B. Value • Less than 1 because it is a weak acid (more reactant compared to product) • Ka increases as the strength of acid increases • Kb increases as strength of base increases Weak acid has Ka < 1 Leads to small [H3O+] Weak base has Kb < 1 Leads to small [OH-] HW # 25, 27

  40. Acids Conjugate Bases Increase strength Increase strength Dr. Mihelcic AP Chemistry 1 HW # 27

  41. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • C. pKa and value • pKa = - log Ka • Value of pKa is smaller as the strength of the acid increases Higher Ka === stronger acid Higher Ka ==== lower pKa

  42. 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. (C3H6O3) Strategy: • Write the equation showing the disassociation of the acid. • Write the expression for Ka • Set up an ICE table and substitute [H+] onto the “E” line. • Calculate the other “E” concentrations. • Solve for the Ka.

  43. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES 2. Write the expression for Ka 5. Solve for the Ka. • Write the equation showing the disassociation of the acid. • I C E 4. Calcother “E” concentrations. • 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. (C3H6O3)

  44. Strategy: • Write the equation showing the disassociation of the acid. • Write the expression for Ka • Set up an ICE table and substitute [H+] onto the “E” line. • Calculate the other “E” concentrations. • Solve for the Ka. • D. Example 17.6 Solution

  45. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES [H+] = pH = % diss = equation • E. Example 17.7 Determine the [H+], pH and % dissociation of carbonic acid (H2CO3 ) in a 0.100 M solution. (Ka = 4.2 x 10-7) I C E Ka=[ ][ ]=[ ][ ]=4.2x10-7 [ ] [ ] X = HW Problem 51

  46. E. Example 17.7 Solution

  47. E. Example 17.7 Including Solving Quadratic Equation:

  48. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • F. Approximation • Small Ka = few protons given off = very little dissociation (acid is a “Relatively Weak” weak acid) • If Ka is smaller than 10-4 • “x” (∆[reactant]) is negligible in “initial - x” term (set x = 0) • For quadratic approxs only! So, for Example 17.7 we could ignore the “x” here: We will still get approx the same value when we solve for x (when you use sig figs…) Yeah- we can avoid solving a quadratic!

  49. Weak Bases

  50. 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 = In our “weak acid” notes: • Ka increases as the strength of acid increases • Kb increases as strength of base increases Weak acid has Ka < 1 Leads to small [H3O+] Weak base has Kb < 1 Leads to small [OH-] • Kb increases as strength of base increases (see exercise 17.5 pg. 705)