Chapter 3
Acids and Bases
 A substance can be classified as an acid or a base depending on
certain properties.
 There are several theories to define and classify acids and bases
that includes:
1. Arrhenius Acid-base theory
2. Bronsted-Lowry Acid Base theory
3. Lewis Acid Base theory
4. Solvent System Acid Base
Every theory of acid and base shows some advantages and some
limitations to overcome; the drawback of each theory.
1. Arrhenius Acids and Bases
An Arrhenius acid is a substance that dissociates
to give hydrogen ions (H+ ) when dissolved in
water.
HCl (g) + H2O
H+(ag) + Cl-(ag)
An Arrhenius base is a substance that dissociates in to
hydroxyl ions (OH- ) when dissolved in water
NaOH + H2O
Na+(ag) + OH-(ag)
 Some typical Arrhenius acids are HCI, HNO3, and HCN,
and some typical bases are NaOH, KOH, and Ba(OH)2.
2
Limitation of the Arrhenius theory
a) Acidic property of salts such as AlCl3 is ignored
b) The theory defined an acid or a base in terms of hydrogen or
hydroxyl compounds only. It does not explain many substances
that are bases with out having hydroxyl (OH-) group
Example. Weak base ammonia (NH3)
c) The theory does not include non-aqueous solvents.
 Acid-base reactions could take place in non-aqueous media
(ammonia, sulfur trioxide)
3
2. Bronsted-Lowry Acids –Bases Theory
In the Bronsted-Lowry definition:
Acid is a proton(hydrogen ion) donor where as base is a
proton (hydrogen ion ) acceptor
 A Bronsted-Lowry acid base reaction results in the transfer of
a proton from an acid to a base
4
Examples
5
Acid-Conjugate base and Base-conjugate acid pairs
According to Lowry and Brønsted, when acid reacts with a
base, there are two conjugate acid base pairs i.e., for every
acid there exists a base which is produced when the acid
loses its proton and for every base there exists an acid which
is produced when the base accepts proton.
An acid is one proton (hydrogen ion, H+) than its conjugate
base and a base is one proton less than its conjugate acid.
H2SO4 and HSO4- are acid-base conjugate pair; and NH4+ and
NH3 are acid- base conjugate pairs.
 The strong acid , the weak its conjugate base and vice verse
6
Molecules or ions that can lose as well as accept proton are called
amphoteric substances. Example: Water
This concept can explain the acidic/basic nature of a substance in
aqueous (H2O) as well as in other protonic solvents like liq. NH3, liq.
HF.
Limitations
This concept cannot explain the acidic character of non-protonic acids
(which cannot give a proton).BF3, AlCl3, etc.
This concept cannot explain a number of acid – base reaction which
take place in the absence of the solvent or without proton transfer.
7
Lux-Flood concept
According to Lux-Flood concept, the base is an oxide donor and the
acid is an oxide acceptor.
Lux-Flood definition is useful for limited systems such as high
temperature reaction involving molten oxides and anhydrous
reaction
8
3. Solvent System Acid Base concept
According to this concept, the solvents usually undergo self ionization
(auto-ionization) and give rise to cations and anions which are called
solvent cations and solvent anions, respectively.
According to this concept :
Acid is defined as a substance which, either by direct dissociation or by
reaction with the solvent gives the cation characteristic of that solvent
 Base is defined as a substance, which either by direct dissociation or by
reaction with the solvent gives the anion characteristic of that solvent.
9
For example, the characteristic cation and anion of H2O are
H3O+ and OH– respectively:
H2O(l)+H2O(l)→H3O+(aq)+OH−(aq)
Thus, all the compounds which produce H3O+ ions in H2O
will act as acids and all the compounds which produce OH–
ions in H2O will act as bases.
2NH3(l) ⇌NH4+ + NH2−
2H2SO4(l) ⇌H3SO4+ + HSO4−
 The solvent system definition also allows for autoionizations which
involve the transfer of an ion other than hydrogen. For example
2BrF3(l) ⇌BrF2+ + BrF4−
10
 HCl is an acid in water since it increases the concentration
of H3O+ when it dissociates:
HCl(aq) + H2O(l) → H3O+(aq) + Cl−(aq)
 HCl also acts as an acid in liquid ammonia since it gives
NH4+ ion on dissociation.
HClacid + NH3(l)→NH4+ + Cl−
 Antimony pentaflouride acts as an acid in liquid BrF3 since it
abstracts a flouride to give BrF2+.
SbF5 acid +BrF3(l) ⇌ SbF-6 +BrF+2
11
In contrast the flouride ion of potassium flouride acts as a
base since it adds to BrF3 to gives BrF4−.
KF base + BrF3(l) ⇌ K+ + BrF−4
 Acid is any substance that increases the concentration of the
solvent cations normally produced by solvent autoionization.
 base is any substance that increases the concentration of the
solvent anions normally produced by solvent autoionization
The definition of acids and bases can be used for both protonic
(e.g. H2O, NH3 etc.) as well as non-protonic solvents
 The definition is applicable for aqueous (H2O) as well as nonaqueous solvents (NH3, HF, H2SO4 etc).
12
• Limitations
Acid base reaction taking place in the absence of a solvent can’t
be explained, i.e., acid-base reaction takes place only in presence
of solvent.
 The concept can’t account for the acid-base reaction occurring
in nonionizing solvents like C6H6, CHCl3 etc.
13
4. Lewis Acids and Bases Theory
According to this theory:
Acid is an electron pair acceptor and Base is electron pair donor.
This concept includes the reactions in which no protons are
involved and the reactions which take place in absence of solvent
Lewis acids are substances which are electron-deficient (or low
electron density) known as electrophiles. Contain vacant orbitals
Lewis bases are substances which are electron-rich (or high
electron density) - known as nucleophiles. Contain lone pair of
electrons
Several categories of substances can be considered as Lewis acids:
positive ions; having less than a full octet in the valence shell,
expandable valence shells.
14
a) Cations : cations are regarded as lewis acids. The strength
of these cations in general, increase with decrease in ionic
radius and increase in the positive charge carried by the
cation
Examples H+, Mn2+, Al3+, Fe2+ are Lewis acids.
(b) Electron Deficient Compounds All molecules which
have a central atom with an incomplete octet act as lewis
acids. For example – BF3, AlCl3 etc.
15
(c) Molecules with Central Atom Containing Vacant dorbitals : All molecules which have central atom with vacant
d-orbitals, can extend their valence. For e.g. SiF4, SnCl4, SbF5,
 Several categories of substances can be considered Lewis bases:
a) Neutral molecules with lone pair of electrons and double bonds
:NH3, H2O, ROH
(b) All anions like
F−, OH−, CN−, etc
.
16
The Hard-Soft Acid-Base Principle (HSAB)
In coordination chemistry, metal ions and ligands have been
seen to have preferential affinity for particular ligands and
metal ions respectively. Depending upon this tendency
Pearson classified metal ions and ligands into hard and soft
acids and bases.
He divided chemical elements into groups according to acidbase (electron pair acceptor –donor) properties of their atoms
or common ions.
a) Hard Acids : These are the metal ions which are small in
size, have high positive charge. These metal ions include H+
,cations from groups 1, 2 and light transition and inner
transition metals.
d-orbitals are either vacant or non-existent
17
(c) Soft Acids: These are the metal ions which are large in size
and usually of low charge
 These are mostly heavy metal ions generally associated with
low (or even zero) positive oxidation state.
a) Hard Bases: these are the anions or neutral molecules which
are having high electronegativity and relatively small size .
Example: O2- , CO32(d) Soft Bases :These are the anions or neutral molecules , donor
atom of easily oxidized , low electronegativity . E.g (
18
19
Principle of Hard and Soft Acids and Bases
According to this principle, hard acids prefer to combine with hard
bases and soft acids prefer to combine with soft bases to form
stable products.
Hard –soft acid base or soft – hard acid base combinations yield
comparatively less stable product
The guiding principle regarding the interaction of electron pair
donors and acceptors is that the most favorable interactions occur
when the acid and base have similar electronic character
20
Some Applications of HSAB Principle:
to explain the relative stability of complexes.
For example , AgI2- is stable, AgF2- but does not exist. Why?
Soft acid-soft base interaction yields stable complex.
used to predict the existence of certain metal ores.
 Example, hard acids like Mg2+, Ca2+ and Al3+ exist as oxides or
carbonates, because oxides and carbonates are hard bases. Where
as soft acids like Ag+, Cu+ and Hg2+ exist as sulphides, because
sulphide is a soft base.
used to predict the feasibility of reaction between two compounds.
Example, Hg(OH)2 dissolves readily in acidic aqueous solution, but
HgS does not. Hg(OH)2 is a product of soft acid-hard base
combination, so it is unstable and dissociates readily. Whereas HgS
is a product of soft acid-soft base combination. Hence it is stable &
does not dissociate easily.
21
Relative strengths of acids and bases
Higher the values of Ka and Kb, stronger the acid or base
Acid strength is the tendency of an acid to donate a proton.
The more readily a compound donates a proton, the stronger an acid
it is.
Acidity is measured by an equilibrium constant.
When a Brønsted-Lowry acid H—A is dissolved in water, an acidbase reaction occurs, and an equilibrium constant can be written for
the reaction
22
 Because the concentration of the solvent H2O is essentially
constant, the equation can be rearranged and a new
equilibrium constant, called the acidity constant, Ka, can be
defined
 It is generally more convenient when describing acid
strength to use “pKa” values than Ka values.
 The larger the Ka value the stronger the acid. The
larger Pka value the weaker the acid!
23
Selected
values
24
Relative Strengths Of Binary Acids
H –X
Electronegativity and size of element effect
The greater the tendency for the transfer of a proton from HX to
H2O, the more the forward reaction is favored and the stronger the
acid
In a periodic group: The strength of binary acids increase from
top to bottom in a group of the periodic table.
Anion radius is directly proportional to acid strength. The larger
the resultant anion’s radius, the stronger is the acid
 negative charge is stabilized when it is spread over a larger
volume.
25
Bond dissociation energy: the weaker the bond, the stronger the
acid.
Bond dissociation energy
569
> 431 > 368 > 297
(kJ/mol)
HF
HCl
HBr
HI
Acid strength Ka
6.6x10-4 < ~106 < ~108 < ~109
Anion radius: the larger the anion’s radius, the stronger the acid.
Anion radius (ppm)
136
< 181 < 195 < 216
(kJ/mol)
HF
HCl
HBr
HI
Acid strength Ka
6.6x10-4 < ~106 < ~108 < ~109
 Size determines acidity down a column. The larger the size of X
the stronger acid
26
In a period:
Across a row of the periodic table, the acidity of H—X
increases as the electronegativity of X increases
The larger the electronegativity difference between H and X,
the more easily the proton is removed and the stronger is the
acid.
The strengths of binary acids increase from left to right
across a period of the periodic table.
Acid strength CH4
<
NH3 <
H2O < HF
Electronegativity :
It is the relative tendency of an atom in a molecule to attract a
shared pair of electrons towards itself
27
Representative Trends in Strengths of Binary Acids
28
Strengths of Oxoacids
As the electronegativity of the central atom (E)
increases the acid strength increases.
Electronegativity
Acid strength Ka
2.5
<
2.8
<
HOI
HOBr
2.3x10-11 < 2.5x10-9 <
3.0
HOCl
2.9x10-8
HClO4 > HClO3 > HClO2 > HOCl
 In HClO4, the presence of maximum number of electronegative
oxygen atoms increases its acidity.
 The other reason for the highest acidity of HClO4 is that the
negative charge after removing hydrogen can distribute over four
oxygen atoms.
29
O
Strengths of Carboxylic Acids
R
C
O
H
Carboxylic acids all have the -COOH group in common;
therefore, differences in acid strength must come from
differences in the R group attached to the carboxyl group.
In general, the more that electronegative atoms are attached in the
R group, the stronger the acid.
I-CH2CH2COOH Cl-CH2CH2COOH CH3-CHClCOOH CH3CCl2COOH
Ka
8.3x10-5 <
1.0x10-4
<
1.4x10-3
<
8.7x10-3
30
Inductive effect (I-effect)
 The strength of acids and bases can be explained by inductive
effect.
 the inductive effect in a molecule is a local change in the electron
density due to electron-withdrawing or electron-donating groups
elsewhere in the molecule
 In case of electronegative groups, due to –I effect, availability of
electrons on the central atom decreased and thus, basicity of a base
decreased.
 Electron donor groups (+I effect) like methyl (-CH3) group,
increase electron density on central atom, thus, increase its basicity
PF3 is weaker base than PH3.
Base strength order NHMe2 > NH2Me > NH3.
31