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