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Ch 16: Acid-Base Equilibria

Ch 16: Acid-Base Equilibria. Brown, LeMay Ch 16 AP Chemistry Monta Vista High School. 16.1: Acids and Bases. * Defined by Svante Arrhenius in 1880’s Arrhenius acids: produce protons; increase [H + ] HCl (aq) → H + (aq) + Cl - (aq) Arrhenius bases: produce hydroxides; increase [OH - ]

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Ch 16: Acid-Base Equilibria

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  1. Ch 16: Acid-Base Equilibria Brown, LeMay Ch 16 AP Chemistry Monta Vista High School

  2. 16.1: Acids and Bases • * Defined by Svante Arrhenius in 1880’s • Arrhenius acids: produce protons; increase [H+] HCl (aq) → H+ (aq) + Cl- (aq) • Arrhenius bases: produce hydroxides; increase [OH-] NaOH (aq) → Na+ (aq) + OH- (aq) or NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)

  3. 16.2: Dissociation of Water • Autoionization of water: H2O (l) ↔ H+ (aq) + OH- (aq) KW = ion-product constant for water • H3O+ (aq) or H+ (aq) = hydronium

  4. 16.3: The pH Scale pX = -log [X] pH = -log [H+] = -log [H3O+] or [H+] = 10-pH pOH = -log [OH-] or [OH-] = 10-pOH [H+][OH-] = KW = 1.0x10-14 -log ([H+][OH-]) = -log KW -log [H+] + -log[OH-] = -log (1.0x10-14) pH + pOH = 14.00

  5. pOH 0 7 14 pH 14 7 0 16.3: The pH Scale • If [H+]<[OH-], then [H+]<1.0x10-7 Ex: pH = -log[1.0x10-10] = 10.00 (basic) • If [H+] = [OH-] Since [H+][OH-] = 1.0x10-14 [H+] = [OH-] = 1.0x10-7 pH = -log[1.0x10-7] = 7.00 (neutral) • If [H+]>[OH-], then [H+]>1.0x10-7 Ex: pH = -log [1.0x10-3] = 3.00 (acidic)

  6. 16.4: Brønsted-Lowry Acids & Bases • Johannes Brønsted (Denmark) Thomas Lowry (England), 1923 • Brønsted-Lowry acids: H+ donor • Brønsted-Lowry bases: H+ acceptor NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq) Base Acid

  7. Conjugated Acid-Base Pairs • For acid “HA”: HA (aq) + H2O (l) ↔ A- (aq) + H3O+ (aq) acid base conjugate base conjugate acid • For base “B”: B (aq) + H2O (l) ↔ HB+ (aq) + OH- (aq) base acid conjugate acid conjugate base

  8. Amphoterism • Amphoteric: capable of acting as either an acid or base H2O (l) ↔ OH- (aq) + H+(aq) Acting as a base Acting as an acid • * Amphiprotic: can accept or donate a p+

  9. Amphoterism Contd. Many transition metal oxides and hydroxides act as amphoteric substances. Most common amphoteric substance is water. Zinc Oxide acting amphoteric: • In acid: ZnO + 2H+ → Zn2+ + H2O • In base: ZnO + H2O + 2OH- → [Zn(OH)4]2- Al(OH)3 acting amphoteric: • Base (neutralizing an acid): Al(OH)3 + 3 HCl → AlCl3 + 3H2O • Acid (neutralizing a base): Al(OH)3 + NaOH → Na[Al(OH)4]

  10. Partner Activity: Write balanced net ionic equations with a partner showing how Al(OH)3 behaves as an acid and a base Ex: Al (OH)3 + OH-  Al (OH) 4- + H2O

  11. Relative Acid-Base Strengths • The stronger an acid (the greater its ability to donate p+), the weaker its conjugate base (the lesser its ability to accept p+). • The stronger a base, the weaker its conjugate acid. • In an acid-base equilibrium, the p+ is transferred from the strongest acid to the strongest base. HSO4- + CO32- ↔ SO42- + HCO3- Stronger acid Stronger base

  12. 16.5: Strong Acids and Bases • Strong acids and bases fully ionize in water (equilibrium is shifted “entirely” toward ions). Strong acids: HI, HBr, HCl, HClO4, HClO3, H2SO4, HNO3 Ex: In 6M HCl solution, 0.004% exist as molecules Strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2

  13. 16.6: Weak Acids • Weak acids partially ionize in water (equilibrium is somewhere between ions and molecules). HA (aq) ↔ A- (aq) + H+ (aq) • Ka = acid-dissociation constant in water • Weak acids generally have Ka < 10-3 • See Appendix D for full listing of Ka values

  14. Ex: Calculate the pH of 2.0 M HCl solution (Ka≈106) • Strong acid, completely dissociated • HCl (aq) → H+ (aq) + Cl- (aq) 2.0 M 0 M 0 M - 2.0 M + 2.0 M + 2.0 M 0 M 2.0 M 2.0 M So: [HCl]initial = [H+]final = [Cl-]final = 2.0 M pH = - log [H+] = - log [2.0] = -0.30

  15. Ex: Calculate pH of 2.0 M HF solution (Ka=7.2x10-4) • Weak acid, partially dissociated • HF (aq) ↔ H+ (aq) + F- (aq) 2.0 M 0 M 0 M - x M + x M + x M (2.0 – x) M x M x M Using quadratic eq’n, 0 = x2 + 7.2 x 10-4x – 1.44 x 10-3 x = 3.7229 x 10-2 or – 3.8669 x 10-2 = [H+] pH = - log [H+] = - log [3.7 x 10-2] = 1.43

  16. Or, since weak acids partially dissociate, assume that [HF]init >> [H+]eq Then, [HF]init – [H+] ≈ [HF]init pH = - log [H+] = - log [3.8 x 10-2] = 1.42 • General rule: if [H+]  5% of [HA], it is better to use quadratic formula.

  17. Percent Ionization of an Acid • Ex: Calculate the % ionization of: • 2.0 M solution of HCl • 2.0 M solution of HF

  18. Polyprotic acids: have more than one H+ to “donate” Ex: H2SO3 (aq) ↔ HSO3- (aq) + H+ (aq) Ka1 = 1st acid-dissociation constant = 1.7 x 10-2 HSO3- (aq) ↔ SO32- (aq) + H+ (aq) Ka2 = 2nd acid-dissociation constant = 6.4 x 10-8 • Ka1>Ka2; 1st H+ dissociates more easily than the 2nd.

  19. * Polyprotic Acids • Ascorbic acid (Vitamin C): • Citric acid:

  20. 16.7: Weak Bases • Partially ionize in water. B (aq) + H2O (l) ↔ BH+ (aq) + OH- (aq) Kb = base-dissociation constant in water In practice, where x = [OH-]

  21. 16.8: Relationship between Ka and Kb Weak base: NH3(aq) + H2O(l) ↔ NH4+(aq)+OH-(aq) Conjugate acid: NH4+(aq) ↔ NH3(aq) + H+(aq)

  22. NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq) + NH4+ (aq) ↔ NH3 (aq) + H+ (aq) H2O (l) ↔ H+ (aq) + OH- (aq) And: Therefore: For a conjugate acid-base pair

  23. In general, when two reactions are added to give a 3rd, the equilibrium constant for the 3rd reaction equals the product of the equilibrium constants of the two added reactions. Furthermore: For a conjugate acid-base pair

  24. 16.9: Salt Solutions as Acids & Bases • Hydrolysis: acid/base reaction of ion with water to produce H+ or OH- • Anion (A-) = a conjugate base A- (aq) + H2O (l) ↔ HA (aq) + OH- (aq) • Cation (B+) = a conjugate acid B+ (aq) + H2O (l) ↔ BOH (aq) + H+ (aq)

  25. Predicting pH of Salt Solutions Consider the relative strengths of the acid and base from which the salt is derived: Ca2+ conjugate acid of strong base Ca(OH)2 NO3- conjugate base of strong acid HNO3 Strong electrolyte Ex: Ca(NO3)2 Neither H+ nor OH- 7

  26. Na+ conjugate acid of strong base, NaOH ClO- conjugate base of weak acid, HClO Weak electrolyte Ex: NaClO OH- > 7 ClO- (aq) + H2O (l) ↔ HClO (aq) + OH- (aq) where x = [OH-]

  27. NH4+ conjugate acid of weak base, NH3 Cl- conjugate base of strong acid, HCl Weak electrolyte Ex: NH4Cl H+ < 7 NH4+ (aq) + H2O (l) ↔ NH3 (aq) + H3O+ (aq) where x = [H+]

  28. 16.10: Acid-Base Behavior & Chemical Structure • Strength of acids and bases explained by proton source and sink • Stability of their anions, factors affecting stability of anions • Stronger acids, HA, have: • H with a higher d+ • Weaker H-A bonds (smaller bond energies) • More stable conjugate bases A- • Stronger oxyacids, HxOz-Y, have: • Central nonmetal “Y” with higher electronegativity • More O atoms Ex: Rank these in order from strongest to weakest: HClO, HClO2, HCl, HBr

  29. 16.11: Lewis Acids & Bases • Lewis acid: “e- pair acceptor” • Brønsted-Lowry acid = H+ donor • Arrhenius acid = produces H+ • Lewis base: “e- pair donor” • B-L base = H+ acceptor • Arrhenius base = produces OH- Ex: NH3 + BF3 → NH3BF3 Lewis base Lewis acidLewis salt 6 CN- + Fe3+ → Fe(CN)63- Lewis base Lewis acidCoordination compound Gilbert N. Lewis(1875 – 1946)

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