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Chapter 16 Acids and Bases

Chapter 16 Acids and Bases. Contents in Chapter 16. 16-1 Arrhenius Theory: A Brief Review 16-2 Brønsted–Lowry Theory of Acids and Bases 16-3 Self-Ionization of Water and the pH Scale 16-4 Strong Acids and Strong Bases 16-5 Weak Acids and Weak Bases 16-6 Polyprotic Acids

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Chapter 16 Acids and Bases

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  1. Chapter 16Acids and Bases

  2. Contents in Chapter 16 16-1 Arrhenius Theory: A Brief Review 16-2 Brønsted–Lowry Theory of Acids and Bases 16-3 Self-Ionization of Water and the pH Scale 16-4 Strong Acids and Strong Bases 16-5 Weak Acids and Weak Bases 16-6 Polyprotic Acids 16-7 Ions as Acids and Bases 16-8 Molecular Structure and Acid–Base Behavior 16-9 Lewis Acids and Bases

  3. H2O NaOH(s) → Na+(aq) + OH-(aq) 16-1 Arrhenius Theory: A Brief Review • Acid: A substance that provides H+ in aqueous solution, e.g., HCl • Base (alkalis): A substance that provides OH– in aqueous solution, e.g., NaOH H2O HCl(g) → H+(aq) + Cl-(aq) Na+(aq) + OH-(aq)+ H+(aq) + Cl-(aq) → H2O(l) + Na+(aq) + Cl-(aq) H+(aq)+ OH-(aq) → H2O(l) • Arrhenius theory did not handle non OH– bases such as ammonia (NH3). • Neutralization reaction: Combination of hydrogen ions (H+) and hydroxide ions (OH–) to form water.

  4. Conjugate base of NH4+ Conjugate Acid of NH3 Conjugate base of H2O Conjugate acid of OH− 16-2 Brønsted–Lowry Theory of Acids and Bases*** • Definition: • Acid: the substance act as H+ donor • Base: the substance act as H+ acceptor

  5. Ionization constants*** • Acid ionization constant CH3COOH + H2O CH3COO– + H3O+ Conjugate acid Conjugate base Conjugate base Conjugate acid • Base ionization constant Conjugate base Conjugate acid NH3 + H2O  NH4+ + OH– Conjugate acid Conjugate base *H2O is an amphiprotic (amphoteric) substance, act as either an acid or a base.

  6. Strengths of conjugate acid-base pairs • The stronger an acid, the weaker its conjugate base. (The stronger a base, the weaker its conjugate acid.) • An acid-base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid-base pair.

  7. Brønsted–Lowry acid–base reaction: weak acid • CH3COOH is only slightly ionized. • Reverse reaction proceeds to a greater extent than does the forward reaction. • H3O+ is a stronger acid than CH3COOH and CH3COO− is a stronger base than H2O.

  8. Brønsted–Lowry acid–base reaction: strong acid • HCl is essentially completely ionized. • The forward reaction proceeds almost to completion. • H3O+ is a weaker acid than HCl and Cl− is a much weaker base than H2O.

  9. More about strengths of acid and bases • Leveling (solvent) effect: The solvent's ability to level the effect of a strong acid (or strong base) dissolved in it. e.g., HI and HBr are leveled to the same acidic strength in H2O. • Differentiating (solvent) effect: The solvent's ability to differentiate the acidic (or basic) strength. Example: • Ka (Kb) values are used to compare the strengths of weak acids (bases).

  10. 16-3 Ionization of Water and the pH Scale • Self ionization of water Ion-product of water, KW: KW = [H3O+][OH–] At 25oC, KW = 1x10–14 KW applies to all aqueous solutions—acids, bases, salts, and nonelectrolytes—not just to pure water.

  11. p function: –log • pH = –log [H3O+] [H3O+] = 10–pH pOH = –log [OH–] [OH–] = 10–pOH • Acidic solution: [H3O+] > [OH–] pH < pOH Basic solution: [H3O+] < [OH–] pH > pOH • Aqueous solution at 25oC: pKW = pH + pOH = 14.00 • Aqueous solution at 25oC: acidic solution: [H3O+] > 1.0×10–7 pH < 7.00 basic solution: [H3O+] < 1.0×10–7 pH > 7.00 neutral solution: [H3O+] = 1.0×10–7 pH = 7.00

  12. The pH scale and pH values of some common materials • pH = −1 ([H3O+] ≈ 10 M) and pH = 15 ([OH−] ≈ 10 M) are possible. • The pH scale is useful only in the range 2 < pH < 12, because the molarities of H3O+ and OH− in concentrated acids and bases may differ from their true activities.

  13. 16-4 Strong Acids and Strong Bases • The contribution due to the self-ionization of water can generally be ignored (unless the solution is extremely dilute), i.e., for strong acids and bases, dissociated completely. Therefore, [H3O+]  CHCl CHCl: initial concentration of HCl [OH–]  CNaOH CNaOH: initial concentration of NaOH.

  14. For extremely dilute solution of a strong acid and strong base, 1.0 x 10–8 M HCl for example:

  15. 16-5 Weak Acids and Weak Bases • Identifying Weak Acids and Bases

  16. [H3O+] from HA Degree of ionization = [HA] originally [H3O+] from HA 100% Percent ionization = [HA] originally • Percent Ionization (A weak acid HA for example) HA + H2O H3O+ + A-

  17. Equilibrium of monoprotic acid, HA HA + H2O  H3O+ + A– CHA– x x x Therefore • 5% rule: Assume CHA – x  CHA * CHA – x  CHA, using:

  18. Equilibrium of monobasic base, B B + H2O  HB+ + OH– CB– x x x Therefore • 5% rule: Assume CB – x  CB * CB – x  CB, using:

  19. 16-6 Polyprotic Acids

  20. Diprotic acid H2A + H2O  HA– + H3O+ HA–+ H2O  A2– + H3O+ For conjugate base: A2–+ H2O  HA– + OH– HA–+ H2O  H2A + OH– Ka1 x Kb2 = Kw Ka2 x Kb1 = Kw

  21. Triprotic acid Phosphoric acid for example: H3PO4 + H2O  H2PO4– + H3O+ H2PO4–+ H2O  HPO42– + H3O+ HPO42–+ H2O  PO43– + H3O+ Ka1 x Kb3 = Kw Ka2 x Kb2 = Kw Ka3 x Kb1 = Kw • Ionization constants for polyprotic acid progressively decrease: Ka1 > Ka2 > Ka3 > ….. • Except in very dilute solutions, essentially all of the H3O+ ions come from the first ionization step alone.

  22. EXAMPLE 16-9 Calculating Ion Concentrations in a Polyprotic Acid Solution For a 3.0 M H3PO4, calculate (a) [H3O+]; (b) [H2PO4–]; (c) [HPO42–]; (d) [PO43–]. Solve (a) assume that all H3O+ the forms in the first ionization step H3PO4 + H2O  H2PO4– + H3O+ Initial conc. 3.0 - - Change –x ＋x ＋x Equilibrium (3.0 – x) x x x2 = 0.021 x = [H3O+] = 0.14 M Check: (x/CH3PO4) x 100% = 4.7% < 5%, OK!!

  23. (b) H3PO4 + H2O  H2PO4– + H3O+ Equilibrium (3.0 – x) x x [H2PO4–]  [H3O+] =0.14 M (c) H2PO4–+ H2O  HPO42– + H3O+ Since [H2PO4–]  [H3O+], therefore [HPO42–] = 6.3 x 10–8 M

  24. (d) HPO42–+ H2O  PO43– + H3O+ Since [H3O+] = 0.14 M and [HPO42–] = 6.3 x 10–8 M [PO43–] = 1.9x 10–19 M

  25. A Somewhat Different Case: H2SO4 • (1)H2SO4 + H2O  HSO4– + H3O+ Ka1 103 • CH2SO4 CH2SO4 CH2SO4 • (2) HSO4–+ H2O  SO42– + H3O+ • CH2SO4－x x x • Concentrated solutions(>0.5 M H2SO4): H3O+ is predominated by first ionization step. e.g., 1.00 M H2SO4, [H3O+]  1.00 M. • Very dilute solutions(< 0.001 M H2SO4): both ionization steps are nearly completely dissociated, e.g., 0.001 M H2SO4, [H3O+]  0.002 M, [SO42–] 0.001 M. • Intermediate concentrations (0.001 M< CH2SO4 <0.5 M), first ionization step is completely dissociated, thesecond ionization step is partially dissociated.

  26. Check: (x/CH2SO4) x 100% = 2.2% < 5%, OK!!

  27. 16-7 Ions as Acids and Bases • Hydrolysis: The reaction between an ion and water. (1) NaCl(aq) Na+(aq) + Cl–(aq) Neutral Na+(aq) + H2O  Cl–(aq) + H2O  (2) NH4Cl(aq) NH4+(aq) + Cl–(aq) Acidic NH4+(aq) + H2O  NH3(aq) + H3O+(aq) Cl–(aq) + H2O  (3) CH3COONa(aq) Na+(aq) + CH3COO–(aq) Basic Na+(aq) + H2O  CH3COO–(aq) + H2O  CH3COOH(aq) + OH–(aq) (4) CH3COONH4(aq) NH4+(aq) + CH3COO– (aq) ????? NH4+(aq) + H2O  NH3(aq) + H3O+(aq) CH3COO–(aq) + H2O  CH3COOH(aq)+ OH–(aq) X X X X

  28. The pH of Salt Solutions

  29. (a) NaOCl(aq) Na+(aq) + OCl–(aq)Basic Na+(aq) + H2O  OCl–(aq) + H2O  HOCl(aq) + OH–(aq) X (b) KCl(aq) Na+(aq) + Cl–(aq)Neutral K+(aq) + H2O  Cl–(aq) + H2O  X X (c) NH4NO3(aq) NH4+(aq) + NO3–(aq)Acidic NH4+(aq) + H2O  NH3(aq) + H3O+(aq) NO3–(aq) + H2O  X

  30. 16-8 Molecular Structure and Acid–Base Behavior • Strengths of Binary Acids • Homolytic dissociation vs. Heterolytic dissociation • Homolytic dissociation • HX  H + X D(H–X) • Heterolytic dissociation • HX  H+ + X– D(H+X– ) • Bond dissociation energy for the gas phase ionization reaction D(H+X– ) = D(H–X) + IE(H) + ΔHea

  31. Bond dissociation energies (kJ/mol ) and Ka values for some binary acids

  32. Comparing binary acids of X in the same row* • The higher polarity of the bond (the larger ΔEN (electronegativity difference)), the stronger acid. • Small ΔEN has more covalent character • Large ΔEN has more ionic character Molecule CH4 NH3 H2O HF ΔEN 0.4 0.9 1.4 1.9 Acidity: CH4 < NH3 < H2O < HF

  33. Comparing binary acids of X the same group The larger bond length (the larger X radius, the weaker H—X bond) the stronger acid. Molecule HF HCl HBr HI BE (kJ/mol) 565 431 364 297 Anion radius (pm) 136 181 195 216 Ka6.6x10–4 ~106 ~108 ~109 Acidity: HF < HCl < HBr < HI *** HF(aq) is a weak acid Other example: Acidity: H2O < H2S < H2Se < H2Te

  34. Strengths of Oxoacids (H–O–EOn) • Comparing the EN of E • The larger EN (electronegativity) of E, the weaker H–O bond, the stronger acid. Molecule H–OI H–OBr H–OCl EN 2.5 2.8 3.0 Ka 2.3x10–11 2.5x10–9 2.9x10–8 Acid strength HOI < HOBr < HOCl More examples: Acid strength: H2SeO3 < H2SO3 HBrO4< HClO4

  35. Comparing the n of On • The more number (n) of terminal O, the weaker protonated H–O bond, the stronger acid. *O is the element has second higher electronegativity. More examples: Ka1(H2SO3)< Ka1(H2SO4)

  36. Strengths of R–COOH vs. R–OH • Ethoxide ion is a much stronger base than is acetate ion. • The stronger the conjugate base, the weaker the corresponding acid.

  37. Strengths of carboxylic acids (R–COOH) • Comparing electron-donating ability of R • Electron-donating ability: –C2H5 > –CH3 > –H • The higher electron-donating ability, the weaker acid strength. Example Acidity: CH3CH2–COOH < CH3–COOH < H–COOH Ka: CH3CH2–COOH < CH3–COOH < H–COOH pKa: CH3CH2–COOH > CH3–COOH > H–COOH

  38. Comparing electron-withdrawing ability of R • Electron withdrawing ability: Cl > Br > I • The higher electron withdrawing ability, the stronger acid strength Example

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