CH4. Acids and Bases
1
Bronsted-Lowry
Bronsted-Lowry definitions:
Acid = proton donor; Base = proton acceptor
HF (aq) + H2O
BL acid
H3O+ (aq) + F- (aq)
BL base
Fluoride ion is the conjugate base of HF
Hydronium ion is the conjugate acid of H2O
2
Amphiprotic species
Amphiprotic – species that can act as BL acid or base
NH3 (aq) +
BL base
H2O 
BL acid
NH4+ (aqu) + OH (aqu)
hydroxide
Kb = base dissociation constant = [NH4+] [OH] / [NH3]
H2O is amphiprotic - it’s a base with HF, but an acid with NH3
3
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
4
BL
acid/base
strengths
5
Kw
Kw = water autodissociation (autoionization) constant
2 H 2O

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
6
Polyprotic acids
Since pKa values are
generally wellseparated, 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
7
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
8
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.
9
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
O2
 OH
NH3(l)
Na (m)
 Na+ (solv) + NH2(solv) + ½ H2 (g)
slow
very strong base
Na+ (solv) + e (solv)
10
Acid/base chemistry of complexes
Aqueous chemistry:
H2O
Fe(NO3)3 
[Fe(OH2)6]3+(aq) + 3 NO3(aq)
2 [Fe(OH2)6]3+ (aq)
Hexaaquairon(III), pKa ~ 3
 [Fe2(OH2)10OH]5+ (aq) + H3O+(aq)
dimer
11
Aqua, hydroxo, oxoacids
aqua acid
M(OH2)xn+
ex: [Cu(OH2)6]2+
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
hexaaquacopper(II) cation
ex: H2CO3 (aq) + H2O  HCO3 (aq) + H3O+(aq)
carbonic acid
CO2 (g) + H2O
bicarbonate
pKa ~ 3.6
12
Trends in acidity
For aqueous ions:
pKa vs z2 / (r++ d)
1. Higher charge is more acidic
pKa of [Fe(OH2)]3+ ~ 3
pKa of [Fe(OH2)]62+ ~ 9
2. Smaller radius is more acidic
Mn2+
Cu2+
early TM
late TM
lower Z*
higher Z*
=> larger radius
=> smaller radius
less acidic
more acidic
Na+ (aqu) = [Na(OH2)6]+
has pKa > 14 so it’s a
spectator ion in aqu soln
13
Anhydrides
Ex:
H2O + SO3 
anhydride
H2SO4
acid form
Acidic
SO3 / H2SO4
“P2O5” / H3PO4
CO2/H2CO3
Basic
Na2O / NaOH
Amphoteric
Al2O3 / Al(OH)3
14
Trends in acidity
15
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!
16
Common acids
H2SO4
(Td)
SO42
Sulfuric acid
Sulfate
H2SO3
(C3v)
SO32
Sulfurous acid
Sulfite
You should know these!
17
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!
18
Pauling’s rules for pKa‘s of oxoacids
1.
Write formula as MOp(OH)q
2.
pKa  8 – 5p
3.
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)
19
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
20
Amphoteric oxides
[Al(OH2)6]3+
Oh

Al2O3 / Al(OH)3  [Al(OH)4]
H3O+
OH
Td
2 [Al(OH2)6]3+(aq) 
pKa ~ 2
[Al2(OH2)10(OH)]5+(aq) + H3O+(aq)
dimer
21
polyoxocations
linear trimer is [Al3(OH2)14(OH)2]7+
Keggin ion
[AlO4(Al(OH)2)12]7+
pH ≈ 4
charge/volume ratios
Al(OH2)63+
3+ / Oh
>
dimer
5+ / 2 Oh
>
trimer
7+ / 3 Oh
--- >
Al(OH)3
neutral
22
Polyoxoanions
VO43(aq)
H3O+

V2O5(s)
orthovanadate (Td)
2 VO43(aq) + H2O  V2O74 (aq) + 2OH (aq)
H3O+
V3O93
V3O105
H3O+
oxo bridge
V4O124
23
Lewis acids and bases
A
LA
+ B: 
LB
A:B
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+
24
LA/LB strengths
LA strength is based on reaction Kf
LA/LB strengths depend on specific acid base combination
Ex:
BCl3 + :NR3 
Kf:
NH3 < MeNH2 < Me2NH < Me3N
BMe3 + :NR3
Kf:
Cl3B:NR3

Me3B:NR3
NH3 < MeNH2 < Me2NH > Me3N
Hrxn 58
74
inductive effect
81
inductive + steric
74 kJ/mol
25
log K and ligand type
26
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)
27
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
28
Descriptive chemistry - Group 13
Expect inductive effect
ex:
BF3 > BCl3 > BBr3 but the opposite is true
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
C6H6
C6H5C(O)R
RC(O)-X: + “AlCl3”  RC(O) + AlCl3X
29
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
30
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
K2MnF6 (s) +
2 SbF5 (l)
KF, H2O2
aqu HF
KMnO4
Sb2O3
transient

“MnF4”
+
2KSbF6 (s)
F transfer

MnF3 + ½ F2 (g)
Dove (1980’s), chemical synthesis of F2 gas
31
Descriptive chemistry - Group 16
Inductive effect stabilizes conjugate base (anionic form)
sulfuric acid
fluorosulfonic
HSO3F / SbF5
pKa ~ 2
pKa ~ 5
pKa ~ 26 (superacid)
C6H6
HSO3F / SbF5

C6H7+ SbF6
32