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CH4. Acids and Bases.

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1 CH4. Acids and Bases

2 Bronsted-Lowry 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

3 Amphiprotic species 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

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 (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  H2PO H3O pKa1 = 2.1 H2PO H2O  HPO H3O pKa1 = 7.4 HPO H2O  PO 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 Solvent leveling 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) NH3(l) slow very strong base  OH O2

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

12 Aqua, hydroxo, oxoacids 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

13 Trends in acidity For aqueous ions: pKa vs z2 / (r++ d)
. 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
“P2O5” / H3PO4 CO2/H2CO3 Basic Na2O / NaOH Amphoteric Al2O3 / Al(OH)3

15 Trends in acidity

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

18 Common acids You should know these! 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  (exptl value is 7.4) pKa3  (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
HlO H2O  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 ~ dimer H3O+ OH

22 polyoxocations linear trimer is [Al3(OH2)14(OH)2]7+
charge/volume ratios Al(OH2) > dimer > trimer > Al(OH)3 3+ / Oh / 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 Lewis acids and bases 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 :CO  [W(CO)6] BCl3 + :OEt2  BCl3:OEt2 D3h Fe3+(g) :OH2 → [Fe(OH2)6]3+

25 LA/LB strengths LA strength is based on reaction Kf
LA/LB strengths depend on specific acid base combination Ex: BCl :NR3  Cl3B:NR3 Kf: NH3 < MeNH2 < Me2NH < Me3N inductive effect BMe :NR3  Me3B:NR3 Kf: NH3 < MeNH2 < Me2NH > Me3N inductive + steric Hrxn    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 H2O 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) SbF5 (l)  “MnF4” KSbF6 (s) F transfer KF, H2O2 aqu HF  KMnO4 Sb2O3 MnF3 + ½ F2 (g) Dove (1980’s), chemical synthesis of F2 gas

32 Descriptive chemistry - Group 16
Inductive effect stabilizes conjugate base (anionic form) sulfuric acid fluorosulfonic HSO3F / SbF5 pKa ~  pKa ~ 5 pKa ~ 26 (superacid) C6H  C6H7+ SbF6 HSO3F / SbF5


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