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CH4. Acids and Bases

CH4. Acids and Bases. Bronsted-Lowry definitions: Acid = proton donor; Base = proton acceptor HF (aq) + H 2 O H 3 O + (aq) + F - (aq) BL acid BL base Fluoride ion is the conjugate base of HF Hydronium ion is the conjugate acid of H 2 O. Bronsted-Lowry.

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CH4. Acids and Bases

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  1. CH4. Acids and Bases

  2. Bronsted-Lowry definitions: Acid = proton donor; Base = proton acceptor HF (aq) + H2O H3O+ (aq)+ F- (aq) BL acid BL base Fluoride ion is the conjugate base of HF Hydronium ion is the conjugate acid of H2O Bronsted-Lowry

  3. Amphiprotic – species that can act as BL acid or base • NH3 (aq) + H2O  NH4+ (aqu) + OH (aqu) • BL acid hydroxide • Kb = base dissociation constant = [NH4+] [OH] / [NH3] • H2O is amphiprotic - it’s a base with HF, but an acid with NH3 BL base Amphiprotic species

  4. BL acid/base strength Ka, the acidity constant, measures acid strength as: Ka = [H3O+] [A-] / [HA] pKa = - log Ka When pH = pKa, then [HA] = [A-] For strong acids pKa < 0 pKa(HCl) ≈ -7

  5. BL acid/base strengths

  6. Kw Kw = water autodissociation (autoionization) constant 2 H2O  H3O+ (aqu)+ OH- (aqu) Kw = [H3O+] [OH-] = 1 x 10-14 (at 25°C) Using the above, you should prove that for any conjugate acid-base pair: pKa + pKb = pKw = 14

  7. Polyprotic acids Since pKa values are generally well-separated, only 1 or 2 species will be present at significant concentration at any pH H3PO4 + H2O  H2PO4- + H3O+ pKa1 = 2.1 H2PO4- + H2O  HPO42- + H3O+ pKa1 = 7.4 HPO42- + H2O  PO43- + H3O+ pKa1 = 12.7

  8. Solvent leveling The strongest acid possible in aqueous solution is H3O+ Ex: HCl + H2O  H3O+ (aq) + Cl- (aq) there is no appreciable equilibrium, this reaction goes quantitatively; the acid form of HCl does not exist in aqueous solution Ex: KNH2 + H2O  K+ (aq) + OH- (aq) + NH3 (aq) this is solvent leveling, the stable acid and base species are the BL acid-base pair of the solvent NH2- = imide anion NR2- , some substituted imide ions are less basic and can exist in aq soln

  9. Solvent leveling Only species with 0 < pKa < 14 can exist in aqueous solutions. The acid/base range for water stability pKw, i.e. 14 orders of mag in [H+]. Other solvents have different windows and different leveling effects.

  10. 2EtOH  EtOH2+(solv) + EtO (solv) K ~ 1020 • chemistry in the range of -3 < pKa < 17 • NH3 NH4+(solv) + NH2(solv) • ammonium imide • chemistry in the range of 10 < pKa < 38 • Na (m)  Na+ (solv) + NH2(solv) + ½ H2 (g) • Na+ (solv) + e (solv) • O2 •  OH NH3(l) slow very strong base Solvent leveling

  11. Aqueous chemistry: • Fe(NO3)3 [Fe(OH2)6]3+(aq) + 3 NO3(aq) • 2 [Fe(OH2)6]3+ (aq)  [Fe2(OH2)10OH]5+ (aq) + H3O+(aq) • Hexaaquairon(III), pKa ~ 3 H2O Acid/base chemistry of complexes dimer

  12. aqua acid M(OH2)xn+ ex: [Cu(OH2)6]2+ hexaaquacopper(II) cation • hydroxoacid M(OH)x ex: B(OH)3 , Si(OH)4 pKa ~ 10 • oxoacid MOp(OH)q p and q designate oxo and hydroxo ligands • ex: H2CO3 (aq) + H2O  HCO3 (aq) + H3O+(aq) • carbonic acid bicarbonate • pKa ~ 3.6 CO2 (g) + H2O Aqua, hydroxo, oxoacids

  13. Trends in acidity • For aqueous ions: • . Higher charge is more acidic • pKa of [Fe(OH2)]3+ ~ 3 • pKa of [Fe(OH2)]62+ ~ 9 • . Smaller radius is more acidic • Mn2+ Cu2+ • early TM late TM • lower Z* higher Z* • => larger radius => smaller radius • less acidic more acidic pKa vs z2 / (r++ d) • Na+ (aqu) = [Na(OH2)6]+ has pKa > 14 so it’s a spectator ion in aqu soln

  14. Anhydrides • Ex: H2O + SO3 H2SO4 • anhydride acid form • Acidic • SO3 / H2SO4 • “P2O5” / H3PO4 • CO2/H2CO3 • Basic • Na2O / NaOH • Amphoteric • Al2O3 / Al(OH)3

  15. Trends in acidity

  16. Common acids • HNO3 NO3(D3h) • Nitric acid Nitrate • HNO2 NO2 (C2v) • Nitrous acid Nitrite • H3PO4 PO43 (Td) • Phosphoric acid Phosphate • H3PO3 HPO32 (C3v) • Phosphorous acid Phosphite You should know these!

  17. Common acids • H2SO4 SO42 (Td) • Sulfuric acid Sulfate • H2SO3 SO32 (C3v) • Sulfurous acid Sulfite You should know these!

  18. Common acids • HClO4 ClO4 (Td) • Perchloric acid Perchlorate • HClO3 ClO3 (C3v) • Chloric acid Chlorate • HClO2 ClO2 (C2v) • Chlorous acid Chlorite • HOCl OCl • Hypochlorous acid Hypochlorite You should know these!

  19. Pauling’s rules for pKa‘s of oxoacids • Write formula as MOp(OH)q • pKa 8 – 5p • Each succeeding deprotonation increases the pKa by 5 • Ex: rewrite HNO3 as NO2(OH) • p = 2; pKa 8 – 5(2)  2 (exptl value is 1.4) • Ex: rewrite H3PO4 as PO(OH)3 • p = 1; pKa1 8 – 5(1)  3 (exptl value is 2.1) • pKa2 8 (exptl value is 7.4) • pKa3 13 (exptl value is 12.7)

  20. pKa values • p Pauling pKa • calcn exptl • Cl(OH) 0 8 7.5 • ClO(OH) 1 3 2.0 • ClO2(OH) 2 2 1.2 • ClO3(OH) 3 7 ≈ 10 HlO4 + 2H2O  H5IO6

  21. Amphoteric oxides • [Al(OH2)6]3+ Al2O3 / Al(OH)3  [Al(OH)4] • Oh Td • 2 [Al(OH2)6]3+(aq)  [Al2(OH2)10(OH)]5+(aq) + H3O+(aq) • pKa ~ 2 dimer H3O+ OH

  22. polyoxocations • linear trimer is [Al3(OH2)14(OH)2]7+ • charge/volume ratios • Al(OH2)63+ > dimer > trimer --- > Al(OH)3 • 3+ / Oh 5+ / 2 Oh 7+ / 3 Oh neutral Keggin ion [AlO4(Al(OH)2)12]7+ pH ≈ 4

  23. Polyoxoanions H3O+ • VO43(aq)  V2O5(s) • orthovanadate (Td) • 2 VO43(aq) + H2O  V2O74 (aq) + 2OH (aq) • V3O93 V3O105 • V4O124 H3O+ H3O+ • oxo bridge

  24. A + B:  A:B • LA LB complex • LA = electron pair acceptor; LB = electron pair donor • Lewis definition is more general than BL definition, does not require aqueous or protic solvent • Ex: W + 6 :CO  [W(CO)6] • BCl3 + :OEt2 BCl3:OEt2 • D3h • Fe3+(g) + 6 :OH2 → [Fe(OH2)6]3+  Lewis acids and bases

  25. LA/LB strengths • LA strength is based on reaction Kf • LA/LB strengths depend on specific acid base combination • Ex: BCl3 + :NR3  Cl3B:NR3 • Kf: NH3 < MeNH2 < Me2NH < Me3N inductive effect • BMe3 + :NR3 Me3B:NR3 • Kf: NH3 < MeNH2 < Me2NH > Me3N inductive + steric • Hrxn58 74 81 74 kJ/mol

  26. log K and ligand type

  27. Drago-Wayland equation • A (g) + :B (g)  A:B (g) • Gas phase reactions (omits solvation effects) • -Hrxn = EA EB + CA CB • look up E, C values for reactants (Table 4.4)

  28. Donor/Acceptor numbers • Commonly used to choose appropriate solvents (Table 4.5) • Donor Number (DN) is derived from Hrxn (SbCl5 + :B  Cl5Sb:B) • higher DN corresponds to stronger LB • Acceptor Number (AN) is derived from stability of Et3P=O:A complex • higher AN corresponds to stronger LA • Ex: THF (tetrahydrofuran) C4H8O • DN AN εdielectric constant • THF 20 8 7 • H2O 18 55 82 • Some Li+ salts and BF3 have similar solubilities in THF, H2O • NH3 is much more soluble in H2O • Most salts are much more soluble in H2O

  29. Descriptive chemistry - Group 13 • Expect inductive effect BF3 > BCl3 > BBr3 but the opposite is true • ex: BF3 is stable in H2O, R2O (ethers) • BCl3 rapidly hydrolyzes due to nucleophilic attack of :OH2 • the lower acidity of BF3 is due to unusually favorable B–X bonding in the planar conformation due to  interaction • “AlCl3“ is a dimer (Al2Cl6) • General trend  larger central atom, tends to have higher CN • Al2Me6 is isostructural with Al2Cl6 • Friedel-Crafts • RC(O)-X: + “AlCl3”  RC(O) + AlCl3X C6H6 C6H5C(O)R

  30. Descriptive chemistry - Group 14 CX4 is not a Lewis Acid Acidity SiF4 > SiCl4 > SiBr4 > SiI4 (inductive effect) ex: 2KF(s) + SiF4(g)  K2SiF6(s) LB LA SiF62 Oh SnF4 and PbF4 have Oh not Td coordination (heavier congener, higher CN) each M has 2 unique axial F and 4 shared F

  31. Descriptive chemistry - Group 15 MF5 does not exist for nitrogen; it’s trigonal bipyramidal for M = P, As SbF5: Sb has Oh coordination (oligomerizes to Sb4F20 or Sb6F30) LB LA transient K2MnF6 (s) + 2 SbF5 (l)  “MnF4” + 2KSbF6 (s) F transfer KF, H2O2aqu HF  KMnO4 Sb2O3 MnF3 + ½ F2 (g) Dove (1980’s), chemical synthesis of F2 gas

  32. Inductive effect stabilizes conjugate base (anionic form) sulfuric acid fluorosulfonic HSO3F / SbF5 pKa ~ 2 pKa ~ 5 pKa ~ 26 (superacid) C6H6  C6H7+ SbF6 Descriptive chemistry - Group 16 HSO3F / SbF5

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