1
Acids
Andreas Libau (Libavius)
(1540-1616)
Gerber – prepared HNO3, HCl and aqua regia (from H2SO4
and salts, isolated citric, acetic, and tartaric acids
Libavius - prepared HCl and aqua regia (dissolves Au)
The world highest tonnage per annum chemical commodity:
H2SO4
Gerber - Jabir ibn Hayyan
(721-815)
2
Acids
R. Boyle (1627 - 1691): changes of color of litmus
L. Lavoisier (1743-1794) : Oxygenium = acid forming
Oxides of nonmetals reacts with water to acids
H. Davy (1779-1829)
J. Liebig (1803-1873)
Hydrogen is released during reactions of acids with
metals = H governs acidic properties
3
Arrhenius Theory of Acids and Bases
Svante Arrhenius
(1859-1927)
Acids:
Taste acidic
Release H+ in aqueous soln.
Reacts with base metals with release of H2
K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb
Color litmus red (R. Boyle)
Neutralize bases
Bases:
Taste bitter
Release OH− in aqueous soln.
Color litmus blue
Neutralize acids
4
Strong and Weak Arrhenius Acids
Strong Acids : completely ionized in water
(100% dissociated)
HNO3 → H+
(aq) + NO3
-
(aq)
HCl, HNO3, H2SO4, HClO4, HI, HBr, HClO3, HBrO3, .....
Weak Acids : partially ionized in water
(0% < dissoc. degree <100%)
CH3CO2H H+
(aq) + CH3CO2
-
(aq)
Organic acids, H2CO3, H3BO3, H3PO4, H2S, H2SO3, ...
5
Strong Arrhenius Bases
Strong Bases : completely ionized in water
(100% dissociated)
KOH(s) → K+
(aq) + OH-
(aq)
Alkali metal hydroxides, alkaline earth metal hydroxides,
soluble hydroxides
No weak bases in Arrhenius theory
6
Brønsted–Lowry Acids and Bases
Wider definition, not limited to aqueous solutions
Acid = proton donor = Arrhenius acid
Base = proton acceptor
7
Proton H+
H+ hydronium H3O+ oxonium
Grotthus mechanism
– extremely fast H+ movement
H9O4
+ = [H3O(H2O)3]+ M. Eigen
lifetime H3O+ 1 - 4 ps
[H3O(H2O)20]+
Protonation of solvent (S) is exothermic
H+ + n S → H(S)n
+ ΔH < 0
8
High Proton H+ Mobility
[H3O]+
9
Brønsted–Lowry Strong and Weak Acids
HSO4
- (aq) + H2O H3O+(aq) + SO4
2-(aq)
Equilibrium constant of proton dissociation
= ionization constant of acid
= acid dissociation constant
KC = [H3O+] [SO4
2-] / [HSO4
-] [H2O] [H2O] ~ 55.6 M
Ka = [H3O+] [SO4
2-] / [HSO4
-]
[ ][ ]
[ ]−
−+
=
4
2
43 .
HSO
SOOH
Ka
pKa = −log Ka
Water is a reagent and solvent
10
Acid Strength
pKa = −log Ka
−7HCl
3.2HF
−1.75H3O+
15.57H2O
12.4CF3CH2OH
10.0PhOH
9.2NH4
+
9.1HCN
4.75CH3COOH
−10HClO4
15 – 18ROH
35NH3
pKaAcid
HA (aq) + H2O H3O+(aq) + A-(aq)
Acid
Strength
increases
Strong acids
pKa negative
[ ][ ]
[ ]HA
AOH
Ka
−+
=
.3
Weak acids
pKa positive
NH3 + H2O H3O+(aq) + NH2
-(aq)
11
Acid Strength and ΔG
ΔG = −RT ln Ka = 2.3 RT pKa
HA (aq) + H2O H3O+(aq) + A-(aq)
Ka = [H3O+] [A-] / [HA]
pKa = −log Ka
> 1
< 1
Ka ΔGpKa
< 0< 0Strong acids
> 0> 0Weak acids
[ ][ ]
[ ]HA
AOH
Ka
−+
=
.3
12
Brønsted–Lowry Strong Bases
NH3 + H2O H3O+ + NH2
pKa
(NH3) > pKa(H2O)
CaO + H2O → Ca(OH)2
13
Brønsted–Lowry Weak Bases
NH3(aq) + H2O NH4
+(aq) + OH− (aq)
K = [NH4
+] [OH- ] / [NH3] [H2O]
Kb = [NH4
+] [OH- ] / [NH3] base dissociation constant
pKb = −log Kb
Equilibrium constant of base protonation by water
= ionization constant of base
= base dissociation constant
14
Brønsted–Lowry Weak Bases
Methylamin
CH3NH2 + H2O CH3NH3
+ + OH-
4
23
33
104.4
][
]][[ −
−+
== x
NHCH
OHNHCH
Kb
pKb = −logKb
15
Lewis Acids and Bases
R
N
R
R
+
F
B
FF R
N
R
R F
B
F
F
Lewisovabaze Lewisovakyselina
Lewis Acids – acceptors of electron pair
Lewis Bases – donors of electron pair
Lewis AcidLewis Base
16
Conjugated Pairs of Acids and Bases
H2O(l) + H2SO4(aq) H3O+(aq) + HSO4
-(aq)
Strong Acid Weak base
Acid Conjugated Base
Base Conjugated Acid
Strong Base Weak Acid
Conjugated pairs of acids and bases are connected by proton
exchange
17
Acid-Base Properties of Water
Water is amphoteric – behaves
as an acid and a base
2 H2O(l) H3O+(aq) + OH-(aq)
Autoionization
Water is a very
weak electrolyte
H2O(l) H+(aq) + OH-(aq)
KC = [H+][OH-] / [H2O]2
Kw = [H+][OH-] Kw = ion product of water
Kw = 1 10-14 M2 [H+] = [OH-] = 1 10-7 M
pKw = 14 in pure water at 25 °C and 101.325 kPa
18
pH and pOH Scales
pH = − log [H+]
[OH-] = 1 10-7 M
pOH = − log [1 10-7 ]
pOH = 7
[H+] = 1 10-7 M
pH = − log [1 10-7 ]
pH = 7
in pure water
[H+][OH-] = Kw
pH + pOH = pKW = 14.00
Constant inaqueous solutions
(ionic product of water)
pH < 7 Solution is acidic
pH = 7 Solution is neutral
pH > 7 Solution is basic
1909
S. P. L. Sørensen
Beer brewing
19
pH and pOH Scales
pH pOH [H+] M [OH-] M
0 14 1.0 10-14
2 12 0.01 10-12
4 10 0.0001 10-10
6 8 10-6 10-8
8 6 10-8 10-6
10 4 10-10 0.0001
12 2 10-12 0.01
14 0 10-14 1.0
20
pH and pOH Scales
[H+] [OH-]
pH
Solution is acidic Solution is basic
10-14
Concentration,M
21
Neutralization
H+ + OH- → H2O ΔH = − 56.9 kJ mol−1
H2SO4 + 2 KOH → 2 H2O + K2SO4
H3PO4 + 3 NaOH → 3 H2O + Na3PO4
2 HCl + Mg(OH)2 → 2 H2O + MgCl2
HCl + NaHCO3 → H2O + NaCl + CO2
H3O+(aq) + OH-(aq) → 2 H2O(l) k = 1.4 1011
2 H2O(l) → H3O+(aq) + OH-(aq) k = 2.5 10−5
22
Ka and Kb of Conjugated Pairs
NH3(aq) + H2O NH4
+ (aq) + OH- (aq)
Kb = [NH4
+] [OH- ] / [NH3] basicity constant of NH3
NH4
+(aq) + H2O H3O+ (aq) + NH3(aq)
Kw = [H+][OH-] ion product of water
Ka = [H3O+] [NH3] / [NH4
+] acidity constant of NH4
+
Ka × Kb = [H3O+] [NH3] / [NH4
+] × [NH4
+] [OH- ] / [NH3] = Kw
Ka × Kb = Kw
23
Ka and Kb of Conjugated Pairs
]][[
][
]][[
][
]][[
23
33
33
23
+−
−+
+
+
=
=
=
HOHK
NHCH
OHNHCH
K
NHCH
NHCHH
K
w
b
a
Ka and Kb relationship
CH3NH3
+ CH3NH2 + H+
CH3NH2 + H2O CH3NH3
+ + OHH2O
H+ + OHKa
× Kb = Kw
pKa + pKb = 14
24
Ka and Kw of Water
Autoionization 2 H2O(l) H3O+(aq) + OH-(aq)
KC = [H3O+] [OH-] / [H2O]2
Kw = [H3O+] [OH-] ionic product
pKw = 14
Ka = [H3O+] [OH-] / [H2O] = Kw / [H2O] acidity constant
pKa = 15.74 [H2O] = 55.6 mol l−1
Water is a weak acid
25
Ka of Oxonium Cation H3O+(aq)
H3O+(aq) + H2O(l) H2O(l) + H3O+(aq)
KC = [H3O+] [H2O] / [H3O+] [H2O]
Ka = [H3O+] [H2O] / [H3O+] = [H2O] = 55.6 mol l−1
pKa = − 1.75
Oxonium cation is a strong acid
26
pKa of Acids
Acidstrengthincr.
HA (aq) H+(aq) + A-(aq)
[ ][ ]
[ ]HA
AOH
Ka
−+
=
.3
pKa = −log Ka
pKa
27
35NH2
−NH3
15 – 18RO−ROH
15.57HO−H2O
12.4CF3CH2O−CF3CH2OH
10.0PhO−PhOH
9.2NH3NH4
+
9.1CN−HCN
4.75CH3COO−CH3COOH
3.2F −HF
−1.75H2OH3O+
−7Cl −HCl
−10ClO4
−HClO4
pKaConjug. baseAcid
Acidstrengthincr.
Conjug.
base
strength
incr.
Ka × Kb = Kw pKa + pKb = 14
28
Reaction Equilibrium
Reactions are shifted to weak acids and weak bases
Strong acids expel weak acids
Strong bases expel weak bases
PhOH + F − PhO− + HF
Weak acid. Weak base Strong base Strong acid
Equilibrium shifted to left
CH3COOH + NH3 CH3COO− + NH4
+
Equilibrium shifted to right
29
Nivelization Effect of Water
Acids stronger
than H3O+ are
completely
dissociated in
water
Basea stronger than
HO− are completely
protonated in water35NH2
−NH3
15 – 18RO−ROH
15.57HO−H2O
12.4CF3CH2O−CF3CH2OH
10.0PhO−PhOH
9.2NH3NH4
+
9.1CN−HCN
4.75CH3COO−CH3COOH
3.2F −HF
−1.75H2OH3O+
−7Cl −HCl
−10ClO4
−HClO4
pKaConjug. baseAcid
30
31
pH of Strong Acids and Bases
0.001 M HNO3 → H+ + NO3
[H+]
= 0.001 pH = − log[0.001] = 3
0.1 M KOH → K+ + OH[OH-]
= 0.1 pOH = − log[0.1] = 1 pH = 14 − pOH = 13
1 10-9 M HCl → H+ + Cl[H+]
= 1x10-9 pH = -log(1x10-9) = 9
H2O → H+ + OH- [H+] = 1 10-7
32
pH of Weak Acids
0
0
2
00
][][
][
][
][][][][
][][
)log(
][
]][[
HAKH
HA
H
K
cHAHHAHA
AH
KpK
HA
AH
K
a
a
HA
aaa
=
=
=≈−=
=
−==
+
+
+
−+
−+
HA H+ + ApH
= ½ pKa − ½ log
cHA
Starting acid
concentration
[H+] is very small in
weak acids,
then [HA] = [HA]0
HA H+ + A-
33
HA H+ + A- )log(
][
]][[
aaa KpK
HA
AH
K −==
−+
xx0.1 – xEquilibrium
xx– xChange
000.1Initial
HCOO-H+HCOOH
3
52
4
1012.4
107.1
1.0]1.0[107.1
]1.0[
]][[
−
−
−
=
=
≈−=
−
=
xx
xx
xx
x
xx
Ka
4.12 10-34.12 10-30.1
pH = − log(4.12 10-3) = 2.39
Solution of 0.1 M HCOOH, Ka = 1.74 10-4 . pH = ?
HCOOH H+ + HCOOx
very small,
then [HA] = [HA]0
pH = ½ pKa − ½ log cHA
34
xx0.01 – xEquilibrium
xx– xChange
000.01Initial
HCOO-H+HCOOH
3
6442
642
4
1022.1
2
)1076.1(4107.1107.1
2
4
1076.1107.1
107.1
]01.0[
]][[
−
−−−
−−
−
=
−−±−
=
−±−
=
−+
=
−
=
xx
xxx
a
acbb
x
xxxx
x
x
xx
Ka
1.22 10-31.22 10-30.01
pH = − log(1.22 10-3) = 2.9
Solution of 0.1 M HCOOH
When x comparable to cHA
Then quadratic equation
35
Factors Influencing Acid Strength
Bond dissociation energy, D(HA)
H-A → H· + ·A
Ionization energy of H, IE(H)
H· → H+ + e−
Electron affinity A, EA(A)
·A + e− → A–
Proton affinity A− AP(A)
H+ + A– → HA
Hydratation (solvation) enthalpy
H+ + n H2O → H+ (aq)
A– + n H2O → A− (aq)
HA
D(HA)
H.
+ A.
IE(H)
H+
+ e + A.
EA(A)
H+
+ A-
AP(A-
)
H+
(aq) + A-
(aq)
ΔH solvatΔHr
HA + H2O H+ (aq) + A- (aq) ΔHr
36
Factors Influencing Acid Strength
C C
O
OH
H
H
C C
O
OCl
Cl
Cl
H+
Me C
Me
Me
O
tBuOH pKa = 16
Weak acid
Negative charge on O
= attracts H+
CH3COOH pKa = 4.75
CCl3COOH pKa = 0.52
Strong acid
= low attraction to H+
37
Factors Influencing Acid Strength
H C
H
H
S
O
OH
O
F C
F
F
S
O
OH
O
pKa = - 2.5 pKa = - 12
38
Acidity along Periods Increases
Hydrides = compounds of H with elements
The higher the electronegativity, the higher the acidity,
the better stabilization of negative charge
NaH = basic hydride:
Na+ (aq) + H- (aq) + H2O (l) → H2 (g) + Na+ (aq) + OH- (aq)
HCl = acidic hydride:
H+ (aq) + Cl- (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)
3.215.73555pKa
H-FH-OHH-NH2H-CH3Hydride
Acidity Increases
39
Acidity along Groups Increases
-RSeH−9HBr
7RTeH−9.5HI
10RSH−7HCl
15-16ROH3.2HF
pKaHXpKaHX
Acidity
Increases
Bond
strength
incr
40
Oxyacids
Oxyacids = OH groups bound to an electronegative
central atom
Strong−32
Very strong−83
Weak21
Very weak70
StrengthpKan
(O)n Y (OH)m
Acidity incr
41
Oxyacids
(O)n Y (OH)m
•Stabilization of anions by mezomeric effekt (rezonance)
•Increasing charge on Y
oxidation number of central atom
•Lowering of charge density on O
electronegativity of central atom
•Removal of e. density by other O atoms (-I)
Facilitate dissociation of H+ = incr acid strength
Y O
O
Y
O O
Y
O
O
O
Y
O
O
O
Acidity incr
42
Oxidation Number of Central Atom
−107O3Cl-O-HPerchloric
−35O2Cl-O-HChloric
7.531Cl-O-HHypochloric
2.003O Cl-O-HChlorous
pKaOx.no.FormulaAcid
43
Electronegativity of Central Atom
10.642.7I-O-HHypoiodic
8.693.0Br-O-HHypobromic
7.533.2Cl-O-HHypochloric
pKaElnegE-O-HAcid
44
Oxyacids
M-O ionic bond
= weak acids (= hydroxides)
NaOH Mg(OH)2 Al(OH)3 Si(OH)4 OP(OH)3 O2S(OH)2 O3ClOH
Mn+ - O2- - H
M-O covalent = strong acids
Acidity increases
45
Oxyacids
Factors Influencing Acid Strength
•oxidation number of central atom
•electronegativity of central atom
46
pKa of Bound Water in Metal Complexes
M O
H
H
Acidity incr
Decreasing
ion radius
47
Acidity of Hydrated Cations
0.62.23[Tl(H2O)6]3+
3.92.16[In(H2O)6]3+
2.61.90[Ga(H2O)6]3+
4.91.89[Al(H2O)6]3+
pKaM-O, ÅCation
48
Solutions of Salts
Salts of strong acid – strong base
HCl + KOH H2O + KCl
H+ + Cl- + K+ + OH- H2O + K+ + ClNo
effect on pH
Neutral
Salts of strong acid – weak base - hydrolysis
HNO3 + NH3 NH4
+ + NO3
-
NH4
+ + H2O NH3 + H+ Acidic soln
pH = 7 − ½ pKb − ½ log c
49
Solutions of Salts
NaOH + HF H2O + Na+ + FF-
+ H2O HF + OHBasic
soln
NH4
+ + H2O + F- + H2O HF + OH- + NH3 + H+
Salts of weak acid – strong base - hydrolysis
Salts of weak acid – weak base - hydrolysis
pH = 7 + ½ (pKa − pKb)
pH = 7 + ½ pKa + ½ log c
50
Buffers
1 M acetic acid (HAc) and 1 M sodium acetate (NaAc) in 1
l of solution. NaAc decreases acidity of HAc.
HAc + H2O Ac− + H3O+
Ac− + H2O HAc + HO−
Buffer function = keep constant pH
1. Addition of H+ forms new molecules of HAc
2. Addition of OH− forms new molecules of Ac−
3. pH is constant
51
Buffers
Weak acid and its salt
pH = pKa + log ([A−] / [HA])
Henderson-Hasselbalch eq
Weak base and its salt
pH = pKa + log ([B] / [BH+])
= 14 − pKb + log ([B] / [BH+])
52
Acid-Base Equilibria
1. Ion product of water Kw = [H+][OH-]
2. Dissociation constant Ka = [H3O+] [A-] / [HA]
3. Electroneutrality [H3O+] = [A-] + [OH-]
4. Mass balance [HA]0 = [A-] + [HA]
53
Titration – Controlled Acid-Base Reaction
Equivalence point
= stoichiometric reaction
Titration of strong acid (HCl)
By strong base (KOH)
54
Titration of Weak Acid by Strong Base
pH = ½ pKa − ½ log [HA]
pH = pKa
pH = salt
55
Titration of Weak Acid by Strong Base
HAc + OH- Ac− + H2O
Start of titration : pH = ½ pKa − ½ log [HA]
Half-way
at [A−] = [HA], tj. 50% neutralized (weak acid. + salt):
pH = pKa + log ([A−] / [HA])
pH = pKa
At equivalence, [A−] = 100 [HA]
NEVER [HA] = 0
pH = pKa + log (100 / 1) = pKa + 2
56
Indicators
Acid and its conjugated base have different color
Phenolphtalein
pH = 7.0 8.5 9.4 9.8 12.0
HIn + H2O H3O+ + In-
57
Phenolphtalein
C O
O
C -
O
O
OH O
HO -
O
- 2 H+
+ 2 H+
Colorless Red
58
IndikátoryFFMOlitmus
0
7
14
pH
Color change
pH = pKIN ± 1
59
Indicators
HIn + H2O H3O+ + InKIN
= [H3O+] [In−] / [HIn]
A color change is recognizable if the other form is present
in at least 10% amount
[H3O+] = KIN ([HIn] / [In−] ) from 10:1 to 1:10
pH = pKIN ± 1
60
Titration of Strong Base by Strong Acid
61
Titration of Weak Base by Strong Acid
[B] = [BH+] pH = pKa + log ([B] / [BH+])
= 14 − pKb + log ([B] / [BH+]) = 14 − pKb
Equivalence pH = 14 − pKb + log ([1] / [100]) = 14 − pKb − 2
62
63
HSAB = Hard and Soft Acids and Bases
Hard Acids
Ti4+, Cr3+, Fe3+, Co3+ , H+
Soft Acids
Cu+, Ag+, Hg+, Hg2+, Pd2+, Pt2+
Hard Bases
NH3, NH2R, N2H4, H2O,
OH-, O2-, ROH, RO-,
OR2, CO3
2-, SO4
2-,
OClO3
-, Cl-, F-, NO3
-,
PO4
3-, OCOMe
Soft Bases
H-, R-, C2H2, C6H6,
CN-, CO, SCN-, PR3,
P(OR)3, AsR3, SR2,
SHR, SR-, I-