Module -1 -3 Acid bases and non aqueous solvents

Module -1 Unit 1. Acids, Bases and Non aqueous Solvents
Lecture 1.
B.Sc.III , Dr. S. H. Burungale
Unit 1. Acids, Bases and Non aqueous Solvents [8]
1.1 Introduction to theories of Acids and Bases-Arrhenius concept, Bronsted-Lowry
concept, Lewis Concept, Lux-Flood Concept (definition and examples)
1.2 Hard and Soft Acids and Bases. (HSAB Concept)
1.2.1 Classification of acids and bases as hard, soft and borderline.
1.2.2 Pearson’s HSAB concept.
1.2.3 Acid–Base strength and hardness-softness.
1.2.4 Applications and limitations of HSAB principle.
1.3 Chemistry of Non aqueous Solvents.
1.3.1 Introduction, definition and characteristics of solvents.
1.3.2 Classification of solvents.
1.3.3 Physical properties and Acid-Base reactions in Liquid Ammonia (NH3) and
Liquid Sulphur Dioxide (SO2).
The Arrhenius theory was first introduced by the Swedish scientist Svante Arrhenius in the year 1887. To
conduct electricity, one must have free moving ions. Svante Arrhenius noticed that the solution of acid
conducts electricity by dissolving the substance in the solution, which dissociates into ions. This concept is
well-known these days, but during that time it was controversial. This theory is known as “Electrolytic
dissociation.”
Water is a neutral substance, which does not conduct electricity. By dissolving some substance in water, it
conducts electricity. These substances are called electrolytes and the process is known as “Electrolytic
dissociation.”
Arrhenius Theory of Acid and Base
According to Arrhenius theory, acid is a substance that gives H+
ion on dissolving in the aqueous solution.
It increases the concentration of H+
ions in the solution. The base is a substance that ionizes OH–
ion by
dissolving in the aqueous solution. The concentration of OH- ions is high in the solution.

Arrhenius acids
Arrhenius acid in the aqueous solution increases the concentration of protons or H+
ions. For example,
hydrochloric acid in the water. HCl undergoes dissociation reaction to produce H+
ion and Cl–
ion, as
explained below. The concentration of the H+ ions is increased by forming hydronium ion.
HCl (aq) → H+
(aq) + Cl–
(aq)
HCl (aq) + H2O(l) → H3O+
(aq) + Cl–
(aq)
Other examples of Arrhenius acids are listed below
NHO3(aq) + H2O(l) → H3O+
(aq) + No3
–
In this reaction, nitric acid dissolves in aqueous water to give hydrogen and nitrate ions.
Arrhenius Base
An Arrhenius base is a substrate that increases the concentration of hydroxide ions in the aqueous solution.
The example for Arrhenius base is highly soluble sodium hydroxide compound in water, which dissociates
to give sodium ion and hydroxide ion.
In aqueous solution, NaOH completely dissolves to give hydroxide ion and sodium ion, to increase the
concentration of hydroxide ions.
NaOH(aq) → Na+
(aq) + OH–
(aq)
Some other examples of Arrhenius base are 1st and 2nd group hydroxides, like LiOH and Ba(OH)2.
LiOH +H2O → Li+
+OH-
In the above reaction, lithium hydroxide dissolves in water to give lithium-ion and hydroxide ion.
Examples of an Arrhenius base are listed below
Limitations of Arrhenius theory
The Arrhenius theory is applicable only in aqueous solution; for example, according to the theory, HCl is
an acid in the aqueous solution but not in benzene, even though it donates H+ ion to the benzene. Also,
under Arrhenius’s definition, the solution of sodium amide in liquid ammonia is not alkaline, even though
amide ion deprotonates the ammonia.
Summary
 According to Arrhenius’s theory of acid-base, acids are those which readily dissociate to give the
hydrogen ions in aqueous solution.

 Alkaline species dissolve in aqueous solution to give hydroxide ions.
Bronsted–Lowry Theory
The Bronsted-Lowry theory (Proton theory of acid and base) is an acid-base reaction theory, introduced by
Johannes Nicolaus Bronsted (Danish Chemist) and Thomas Martin Lowry (English Chemist) in 1923.
According to the theory, acid and base react with each other and by an exchange of proton acid, forms its
conjugate base and the base forms its conjugated acid.
The Bronsted-Lowry theory is an extended version of an Arrhenius theory of acid-base.
According to the Arrhenius theory, in aqueous solution, acid increases the concentration of H+
ions and
base increases the concentration of OH–
ions. The limitations of Arrhenius theory was, it identifies the
reaction of an acid and base only in the aqueous medium.
 Strong Bronsted-Lowry acids are those which have a strong tendency to give a proton and their
corresponding conjugate base is weak.
 Weak Bronsted-Lowry acids will have a little tendency to donate a proton and their corresponding
conjugated base is strong.
o When ammonia dissolves in water, hydrogen ions are transferred from water to ammonia
to form ammonium ions and hydroxide ions.
 Ammonia is a Brønsted-Lowry base because it accepts hydrogen ions.

 Water is a Brønsted-Lowry acid because it donates hydrogen ions.
The Bronsted-Lowry acids and their Conjugated Bases
Conjugate Acids and Bases When the temperature of an aqueous solution of ammonia is increased,
ammonia gas is released.
• HNH4
+
reacts with OH–
to form more NH3 and H2O.
• In the reverse reaction, ammonium ions donate hydrogen ions to hydroxide ions.
NH4
+
(the donor) acts as a Brønsted-Lowry acid, and OH−
(the acceptor) acts as a Brønsted-Lowry
base
In essence, the reversible reaction of ammonia and water has two acids and two bases.
• The ammonia molecule and the ammonium ion are a conjugate acid-base pair.
• The water molecule and the hydroxide ion are also a conjugate acid-base pair.
•

In this reaction, hydrogen chloride is the hydrogen-ion donor and is by definition a Brønsted-
Lowry acid. Water is the hydrogen-ion acceptor and a Brønsted-Lowry base The chloride ion is the
conjugate base of the acid HCl.
• The hydronium ion is the conjugate acid of the water base.
.The strength of the acid decreases as it descends and the strength of their corresponding conjugate base
increases.
Summary:
 A Bronsted-Lowry acid is a substance which donates a proton or H+
ion to the other compound and
forms a conjugated base.
 A Bronsted-Lowry base is a substance which accepts a proton or H+
ion from the other compound
and forms conjugated acid.
 Strong acids and bases ionize completely in an aqueous solution, whereas weak acids and bases are
partially ionized in aqueous solution.
 Water molecule is amphoteric in nature, which means it can act as Bronsted-Lowry acid as well as
Bronsted-Lowry base.
Lewis Acids and Bases
In 1923 G. N. Lewis suggested another way of looking at the reaction between H+
and OH-
ions. In the
Brnsted model, the OH-
ion is the active species in this reaction it accepts an H+
ion to form a covalent
bond. In the Lewis model, the H+
ion is the active species it accepts a pair of electrons from the OH-
ion
to form a covalent bond.
In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of
electrons. A Lewis acid is therefore any substance, such as the H+
ion, that can accept a pair of nonbonding
electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such

as the OH-
ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair
donor.
One advantage of the Lewis theory is the way it complements the model of oxidation-reduction reactions.
Oxidation-reduction reactions involve a transfer of electrons from one atom to another, with a net change in
the oxidation number of one or more atoms.
The Lewis theory suggests that acids react with bases to share a pair of electrons, with no change in the
oxidation numbers of any atoms. Many chemical reactions can be sorted into one or the other of these
classes. Either electrons are transferred from one atom to another, or the atoms come together to share a
pair of electrons.
The principal advantage of the Lewis theory is the way it expands the number of acids and therefore the
number of acid-base reactions. In the Lewis theory, an acid is any ion or molecule that can accept a pair of
nonbonding valence electrons. In the preceding section, we concluded that Al3+
ions form bonds to six
water molecules to give a complex ion.
Al3+
(aq) + 6 H2O(l) Al(H2O)6
3+
(aq)
This is an example of a Lewis acid-base reaction. The Lewis structure of water suggests that this molecule
has nonbonding pairs of valence electrons and can therefore act as a Lewis base. The electron configuration
of the Al3+
ion suggests that this ion has empty 3s, 3p, and 3d orbitals that can be used to hold pairs of
nonbonding electrons donated by neighboring water molecules.
Al3+
= [Ne] 3s0
3p0
3d0
Thus, the Al(H2O)6
3+
ion is formed when an Al3+
ion acting as a Lewis acid picks up six pairs of electrons
from neighboring water molecules acting as Lewis bases to give an acid-base complex, or complex ion.
The Lewis acid-base theroy explains why BF3 reacts with ammonia. BF3 is a trigonal-planar molecule
because electrons can be found in only three places in the valence shell of the boron atom. As a result, the
boron atom is sp2
hybridized, which leaves an empty 2pz orbital on the boron atom. BF3 can therefore act as
an electron-pair acceptor, or Lewis acid. It can use the empty 2pz orbital to pick up a pair of nonbonding
electrons from a Lewis base to form a covalent bond. BF3 therefore reacts with Lewis bases such as NH3 to

form acid-base complexes in which all of the atoms have a filled shell of valence electrons, as shown in the
figure below.The Lewis acid-base theory can also be used to explain why nonmetal oxides such as
CO2 dissolve in water to form acids, such as carbonic acid H2CO3.
CO2(g) + H2O(l) H2CO3(aq)
In the course of this reaction, the water molecule acts as an electron-pair donor, or Lewis base. The
electron-pair acceptor is the carbon atom in CO2. When the carbon atom picks up a pair of electrons from
the water molecule, it no longer needs to form double bonds with both of the other oxygen atoms as shown
in the figure below
One of the oxygen atoms in the intermediate formed when water is added to CO2 carries a positive charge;
another carries a negative charge. After an H+
ion has been transferred from one of these oxygen atoms to
the other, all of the oxygen atoms in the compound are electrically neutral. The net result of the reaction
between CO2 and water is therefore carbonic acid, H2CO3.
Module- 2
Hard and Soft Acids and Bases (HSAB). [05]
Classification of acids and bases as hard and soft.
Theoretical bases of hardness and softness
Pearson’s HSAB concept.
Acid – Base strength and hardness and softness.
Application and limitations of HSAB principle.
Introduction :
The Lewis concept failed to provide a definite and uniform scale to measure the relative strength of acids
and bases.The Lewis definition recognizes acid and bases in terms of their ability to accept or donate
electron pairs. The strength of an acid or a base can be determined by the very nature of the reaction
involved in a particular electron transfer process.On the basis of phenomenological criteria, suggested by

Lewis, one may predict that the displacement titrations can be made the basis for much determination.
For example in the reaction :
A+A’B = AB + A’
A’B is converted to AB one may predict that a is stronger than A’. It may be said that the relative
stabilities of acid- base complexes are used to express the relative strengths .From above reaction.AB
must be more stable than A’BAttempts have been made to correlate the different factors governing
strength , from enthalpies (∆H0
) of acid – base reactions .One of the difficulties in such determination
was , with different reference acids (or bases) different trends were observed in Kf , ∆G0
or ∆H0
The complexing ability of the halide ions (Lewis bases)towards Al3+
increases in the order I-
<Br-
< CI-
<F-
But towards Hg2+
the order is just reverse I-
> Br-
>CI-
>F-
.
A similar reversal is seen in the heats of reaction of the acids I2 and C6H5OH with the bases (C2H5)2O
and ((C2H5) 2S. Heat of reaction of I2 is greater with (C2H5) 2S than that with (C2H5) 2O. But the trend for
phenol is just reverse. Inspite of such difficulties ,to deal with the interactions of acids and bases
containing elements drawn from throughout the periodic table ,a qualitative correlation between various
Lewis acids and bases has been achieved .
n 1958 this was done by Ahrland , Chatt and Davies by classifying the acids or bases into general
categories .Class – “a” and Class- “ b” According to them the two categories of metal ions (Lewis
acids)are as follows –
I) Class (a) – The metal ions which prefer to from stable complexes with the ligands having donor atom
of the first members of Gr.15 th(N),16 th(O), and 17
th
(F) in the periodic table . .
Examples are – Alkali metals , Alkaline earth metals and the first row transition metals in high
oxidation state (e.g.fe3+
,Co3+
.Etc. belong to class (a) acids ) .
ii)Class (b )- The other metal ions which prefer to from their most stable complexes with the ligands
having donor atom of the lower members of Gr.15 th(P,As,Sb),16 th(S,Se,Te), and 17 th
(Cl,Br,I) in the
periodic table . .
Lighter transition elements in low oxidation state and heavier transition elements
,say Cu+
, Ag+
,Hg+
,Pt2+
,Pd2+
etc. act as class (b) acids.
CLASSIFICATION OF ACIDS AND BASES AS HARD AND SOFT.

AH:B1+ As+B2=AH :B2+As+B1
From above double displacement reaction it may be stated that B1 is softer than B2 when K1 > 1 .On this
basis a list of hard and soft acids and bases may be obtained .See Table .1.1 The classification is not rigid
and there occurs a gradation from hard acids to soft acids ,and hard bases to soft bases , including the
borderline species. The criterion of hardness (or softness) is ascribed to the “hardness” of the electron
cloud associated with a particular species .A firmly held electron – cloud having low polarizability
makes the species “hard” while an easily polarizable electron clowd characterisesthe species as
“soft”.The third category with intermediate characters will be a borderline .The details of distinguishing
features of hard and soft acids and bases are summarized in table 1.1
Table 1.1 Classification of Lewis Acids and Bases
(A) Acids –
Features :
Hard Borderline Soft
H+,
Li+
,Na+
, K+
Fe
2+
Co
2+
Ni
2+ Cu+
,Ag+
,Au+
,TI+
Be2+,Mg2+,Ca2+ Cu2+
,Zn2+
,Pb2+
Hg+
,Pd2+
,Cd2+
,
Cr
2+
,Cr
3+,
Al
3+ SO2,BBr3 Pt2+
,Hg2+
,BH3,Br2
SO3,BF3,BCl3 Br+
,
HX(H-bonding) M0
(metal atoms) and
bulk metals

(B) Bases :
Hard
Acceptor atoms are marked by:
Soft
Acceptor atoms are marked by:
1) Small size 1) Large size
2) High positive oxidation state. 2) Zero or low positive oxidation state
3) Absence of any outer electrons Which are
easily excited to higher States.
3) Presence of several excitable valence shell
electrons.
4) Absence of d-electrons. 4) With nearly full d-electrons.
5) Usually light metal ions. 5) Mostly heavy metal ions.
6) Know as Lewis acids which Are not easily
polarizable Prefer to coordinate with hard
6) Known as Lewis acids and are easily polarizable.
Prefer to
coordinate with soft bases.
Hard Borderline Soft
F-
,OH-
,H2O,NH3 NO2,SO3,Br-
H-
,R-
,CN-
,CO,I-
,
CO3
2-
,NO3
-
,O2-
, N3
-
,N2,C6H5N, SCN-
,R3P,C6H6,
SO4
2-
, PO4
3-
CIO4
-
,(Cl-
) SCN-
R2S.
Hard Soft
1) High electronegativity. 1) Low electronegativity
2) Low polarisability. 2) High Polarisability.
3) Presence of filled orbits ; empty
orbitals may exist at high energy level.
3) Partially filled orbitals, empty
orbitals are low- lying.
4) These are anions or neutral molecule known as
Lewis bases or ligands, prefer
to co-ordinate with hard acids.
4) these are anions or neutral molecules called
similarly as Lewis bases or
ligands, prefer to bind with soft acids.

Features :From the Table 1.1 it is clear that there is no line of demarcation between hard soft
species.Within each group ,there exists no equal hardness or softness e .g. Alkali metal ions are all hard but
within the group : Li+ >Na+ > K+ > Rb+ >Cs+ hardness decreases hence Li+ is hardest while Cs + is softer
,as it is larger and more polarizable as compared to Li+Similarly nitrogen is a hard base say as NH3 being
of a small size and if polarizable substituents are present, it will turn to be sufficiently softer e.g. Pyridine
,where polarizable substituents are present .Further we may use terms such as “a moderately weak and
fairly soft” “very hard but weak” by considering the strengths of acids and bases.
PEARSON’S CONCEPT
In 1963 R. G. Pearson extended and generalized the qualitative correlation between Lewis acids and
Lewis bases by classifying them into two categories Hard and Soft.
The class –‘a’ metals which are small and less polarizable, prefer to combine with non- metals or ligands
which are also small and not very polarizable , pearson called such metals as Hard Acids and the
corresponding ligands as soft Bases.Similarly the class ‘b’metals having large size ,more or easily
polarisable, prefers to combine with non-metals or ligands having similar properties Pearson called such
metals as soft acids and the ligands as soft baseThe attempt of classification of acidsand bases as hard
and soft by Pearson is known as Hard and soft Acids and Bases .(HSAB) or pearson’s conceptPrinciple
of Pearson’s concept :Pearson suggested a simple rule (Sometimes called Pearson’s principle ) for
predicting the stability of complexes formed between hard and soft acids and bases.
Hard acids prefer to bind (co-ordinate ) with hard bases and soft acids prefer to bind with soft bases and
gives stable complex compound ”.t should be noted that the statement given above is not a theory or an
explanation but it is simple rule of thumb which enables us to predict the relative stabilities of acid-bases
adducts qualitatively.
1;2Theorotical basis of hardness and softness
Several theories have been given to explain the stability of complexes Formed by hard-hard and
soft-soft interactions. Some important theories are:
1. Ionic and covalent bond theory : According to this theory ionic bond is formed by the interaction
of hard acids and hard bases wheras covalent bond is formed by the interaction of soft acids and soft
bases. The electrostatic force of attraction between two oppositely charged ions is inversely
proportional to the internuclear distance. The internuclear distance will be less in case of smaller
ions. Therefore, the electrostatic attraction between two ions will be greater and consequently the

resulting compound will be highly stable.
2.Covalent bond is formed by the interaction of soft acids and soft bases. This is because the soft acids
and soft bases have laege size. The polarization effects are, therefore important to explain their
interactions. Soft acids are generally transition metal ions having six or more d- electrons. The d-sub
shell are easily polarized. Therefore, the complexes formed by soft acids and soft bases have covalent
bonding and are stable.
In order to predict the hard and soft nature of given acid or base, Misons and his coworkers (1967) gave
the following relation-
pK=-logK=AX+BY+C
wher X and Y are the parameters for the acids, A and B are the parameters for the bases, C is a constant
which adjust pK values in such a way that all of them lie on the same scale and K is the equilibrium
constant for the dissociation of acid base complex. The values of parameter Y for some of the acids
(cations) are given below-
The acid
is hard if
the
value of
paramet
er Y is less than 2.80 and the acid is soft if the value of Y more than 3.20. For border line acid the value
of Y is in between 2.80 and 3.20.
e value of parameter B for some of the bases is given below-
Hard base Parameter B Soft base Parameter B
OH-
0.40
I
- 7.17
NH3 1.08
S2O3
2- 12.40
Cl-
2.49
The base is hard if the value of parameter B is less than 3.0 and the base is soft if the value of B is more
then 5.0.
The value of X and A also give information about the hard and soft nature of an acid and a base.
Hard acid Parameter Y Soft base Parameter Y
Li+
0.36 Cu+
3.45
Al3+ 0.70 Tl+
3.78
Na+
0.93 HG2+ 4.25
Ca2+ 1.62 Au+
5.95
Fe3+ 2.37

3 3
3 3 3
3
3 3
1. Bonding theory : This theory was given by Mulliken (1955) and chatt (1956) to explain soft –soft
interaction on the basis of -bonding. Soft acids have low oxidation state and have a large number of d-
electrons. Thus, they have a strong tendency to form -bonds with soft base which are also good –bonding
ligands. The polarization of soft acids and soft base also favour -bonding.
2.Pitzer’s theory : According to Pitzer, London dispersion energies stabilize a bond between two large
polarizable atom . These energies increase with an increase in the size and polarizability . this is
why,soft- soft
interactions are more stable as compared to soft-hard interaction.
ACID-BASE STRENGTH AND HARDNESS-SOFTNESS
Inherent acid –base strength is quite distinguished feature from the hardness and softness.
Hardness - softness pertains to the stability achieved due to hard-hard and soft – soft interactions.
The insight can be collected from the following observations .
i) OH-
and F-
are hard bases where OH-
is 1013 times stronger base than
F
-
ii) Et P and So
2-
are both soft bases where Et P is 107 times stronger base than so2- towards CH
Hg+These facts pertaining to inherent strength violate the Pearson’s principal “ Hard prefers
hard soft prefers soft ”
Soft base SO 2- can displace hard base F-
SO3
2- +HF -------------HSO3
-
+F-
Keq=104
Hard base OH-
can displace soft base SO3
2-
From soft -soft combination of CH3HgSO3
-
OH-
+ CH HgSO3
-
-------------CH HgOH+SO3
2-Keq=10
In these cases the strengths of bases are SO3
2- > F-
and OH-
> SO3
2-are enough to force the
reactions to right irrespective of hard soft considerations.
If both strength and hardness softness are applied under competitive conditions the hard soft rule will be
found to be applicable.
e.g.

3 33
3
3
CH3 HgOH+ HSO3
-
------------CH HgSO3 -
+ HOH Keq 107 soft hard Hard Soft
ii) CH3 HgF +HSO3
-
. -----------CH3HgSO3
-
+HF-
–Keq 103
While acid-base interactions are considered one has to account both
Strength as well as hardness softness.
Tble 1.2: Basicity toward (H+
) and (CH3Hg+
)
Table 1.2 enlists the strengths of different bases toward methylmercury cation CH 3Hg + and the proton
(H+) .From the data it seems that the bases such as triethylphosphine (Et P) and the sulphides S2 –ion are
very strong toward bothCH 3 Hg+ and H+ But both Et3 P and S2-
ion are about a million times better
toward CH3Hg+ , hence both are considered to be soft bases.The OH- ion is a strong base toward both
acids CH3Hg+ and H+ but it is million times better base towards acid,H+
Hence OH- is hard.The F-
ion
is not a good base toward CH3Hg+
or H+
but little better toward H+ as it appears from its hardness .
APPLICATIONS AND LIMITATIONS OF HSAB PRINCIPLE
Applications
With the help of HSAB a large number of chemical reactions can be understood.
Base Linked atom pk(CH3Hg+
) pkn(H+
)
F
- F 1.5 2.58
I
- I 8.6 -9.5
OH-
O 9.37 15.7
S
2- S 21.2 14.2
SO
2-
3
S 8.11 6.79
NH3 N 7.6 9.42
Et3P P 15.o 8.8
CN-
C 14.1 9.14

2 3
4 4
3
Relative strength of Hydracids HF,HCL,HBR and HI :
In aqueous solution the relative strength of HF,HCI,HBr and HI can be predicted.
The reaction of acids with water is:
HX+H O→H O++X-
The hardest base F-
will be most successfully and strongly bonded to the hard acid H+
Hence HF
Will be highly stable.It is therefore least dissocated.Hence acid strength increases as :
HF < HCI < HBr < HI
Relative stabilities of complexes in Aqueous Solutions :
HSAB entails that[ Cd(CN )]2- is more stable that [Cd(NH3) ]2+ According to HSAB principle hard
prefers hard and soft prefers soft.Hence the soft acid Cd2+will prefer to corrdicate soft base CN-
It
is clear from the Kinst constants where
cyano complex has K cyano is stable.
To Predict the Course of Reaction :
H+ CH3HgOH→H2O+CH3Hg+
H+ +CH3HgSH→H2S+CH3Hg+
The reaction (i) goes to right as the hard acid H+
binds strongly to hard base OH-
to produce stable
product H2OOn the other hand the reaction (ii) is favoured to left where soft base SH will tend to
remain combined with soft acid CH Hg+
instead of joining to hard acid H+
Classification of Cations : Fig 1;2
The Fig 1.2 shows the trends in equilibrium constants for formation of complexes with halide ion –
bases.
The Values for Kf increase steeply from F-
to I-
when acid Hg 2+ is used indicating it to be markedly soft
.The curve is less steep but in the same direction for Pb2+ which indicated that Pb2+ is a borderline soft
acid .The trend is in opposite direction for Zn2+ with moderate steepness from I-
to F- .Hence it must be
hard but borderline hard acid .The very steep downward slope for AI3+ indicates clearly that AI 3+
must

be a hard acid .Thus the opposing trends in the reactivity of the halide ions towards Al3+
(F-
> > CI -
> Br
- > I- ) and Hg2+ (F-
> > CI -
> Br -
> I-
) are now easily rationalized .
Classification of Netutral Molecular Compounds :
For neutral molecular acids and bases a similar Classification hard and soft can be applied .For example
the Lewis acid phenol (C6 H5OH) forms a more stable complex by hydrogen bonding to diethyl ether
(C2 H5) 2O than the thioether (C2 H 5 )2 S. In contrast the Lewis acid I2 forms a more stable complex
with (C2 H5)2S Hence we conclude that phenol is hard whereas I2 is soft.
. Pauling Pearson Paradox :
Hard hard and soft soft combinations determine the course of number of typical reactions for
example
LiI + CsF →LiF + CsI HgF2 + BeI2 →BeF2 + HgI2
These reactions will illustrate the Pauling – Pearson Paradox of chemistry .As far as pauling’s electro -
negativity concept is concerned caesium and mercury should form more stable bonds with fluorine as
their electro negativity differences are greater.
In reality however LiF if more stable thatn CsF (and BeF2 than HgF2 ) It is due to very large contribution
of electrostatic interaction between Li+
and F-
and (Be2+
and 2F -
)
In Fact the major driving force for the above combinations comes from the stability of the hard - hard
combination occurring between small atoms joining by ionic bonding and/or covalent bonding .The soft -
soft combinations contribute little or nothing to the driving force except when other factors like
bonding are involved .The following data of enthalpies of atomization will support the facts.
Li I + CsF → Li F + CsI
Hard –soft Hard -soft Hard- soft Hard -soft
347 501 573 335 kjmol-1
Hg F 2 + BeI2 → BeF2 + HgI2
Hard-Soft Hard- Soft Hard- Soft Hard -Soft 535

3
3 5
2 2
6 6 3 5
577 1262 292
kjmol-1
Symbiosis :
BF3 is a hard acid combines readily with a further F- ion which is a hard base .While BH being a soft
acid prefers to join the softer base H-
ion .This fact will easily account for the following :
BF3 H-
+ BH3 F –
→ BF4 -
+ BH3 -
CF3H + CH3 F → CF4 + CH4
Such tendencies of fluoride ions or hydride ions to favour further co ordination by a fourth F-
and H-
ion
has been termed “symbiosis” by Jorgensen (1964) for the symmetrically substituted species with a
centre already having soft ligands or vice Versa.
8. General Chemical Aspcts :
Number of chemical aspects have been interpreted by the HSAB or SHAB concept.
Catalytic power of metals may be accounted from the fact that the soft metal atoms will easily
adsorb soft bases on their surface.
Solubilities may be understood from the fact that the hard solvents will prefer to dissolve hard
solutes and soft solvents dissolve soft solutes e.g .Hg.(OH)2 dissolves in acidified aqueous
solvent but HgS does not.
Substitution reactions can be kinetically studied from hardness and softness of species
concerned.
Some additional illustrations of HSAB concept are
MgCO3 CaCO3 AI2O3 occur in nature but MggS CaS or AI2S3 donot
CU+
Ag+
Hg+
occur in nature as sulphides.
[CoF ]3- is more stable than [CoI ]3-,[Co(NH ) F]2+ is stable [Co(NH ) I]2+ is
unstable.ets.
[Ag(CN) ]-
is very stable but[AgCI ]-
is very unstable.

2
2
f 2 2 3 3
3 2 53 2 5
[AgI ]-
(produced by soft-soft combination)is stable and exists while [AgF ]-
( produced by soft-hard
combination) does not exist.
Ammonia.water.fluoride ion ets. Prefer to bind to Be2+ Ti4+ Co3+ ect and give very stable
complxes.
vii) Phosphines (R3P) thioethers (R2S) and other species of P and S as donor atoms prefer to
bind to Pt2+ Pd2+ Hg2+ ect .
Referring to Table .1.1 all such cases may by easily tackled by soft hard - hard soft - soft or hard soft
combinations relating to the person rule. Energetic of Hardness :In general hard acids are identified
empirically by their preferential binding of lighter basic atoms within the group.
e.g. For hard acids : Kf +F- >> CI-
> Br-
> I-
> R2O> > R2S R3N > > R3P Conversely the soft acids are
indentified empirically by showing the opposite trend down the groups:
For soft acids :K = F-
< < CI-
< Br-
< I -
R O< < R S R N << R P
Limitation of HSAB Concept /Principle
Hard and soft classification is useful concept no doubt but it has some tricky limitations as
pointed out below.
The prime limitation f the HSAB concept is that it is widely general and has no any direct
quantitative scale of acid base strength .
The inherent acid base strengths are not accounted for e.g.OH- and F- ions are both hard bases where
OH- is nearly 1013 times stronger base than F ions .Correlation between hardness and inherent acid
base strength is yet to be developed.
Interpretation of different reactions by splitting the participants into acid base fragment is quite arbitrary
to some extent. The reaction between ethanol and acetic acid may be interpreted for esterification in two
ways:
Break I CH COO-H + CH OH- Break II Ch Co+ OH
+ C H O-H+
The hard-hard combination of H+

3 2 5
The hard combination of H+ with OH- for both is justifiable .But there is nothing to exclude the break
(I) on the basis of hard –soft interactions between CH COO – and C H +
Sometimes Hard Soft principal fails to keep parity with inherent acid-base strengths.
This reaction must be favoured in the view of soft soft combination between CH3 and H- .But in
actual practice the combination is endothermic by about + 360KJ moI-1 .This unfavourable entropy
chage dose not allow the reaction to proceed.
(In favour of reaction the explanation may be given in terms of the greater acidity of proton H+
relative to CH3+ cation )
Hard soft combinations occur in many cases.
e.g. SO2-3 +HF – HSO3-+F
Here it appears that the soft base SO23- has replaced the hard base
F- and combines with hard acid H+ (Here the soft base SO23- must be stronger than the hard base F-)
With these few illustrations it is worth to recall R.G.Pearson who says.
It should be stressed that the HSAB principal is not a theory but is a mere statement about experimental
facts Accordingly an explanation of some observation in terms of hard and soft behavior does not
invalidate some theoretical explanation.
CONCLUSION
To sum up it may be predicted that the importance of HSAB approach lies in its utility to systematize a
good number of experimental observations leading to the study of relative stabilities of complexes. The
point worth nothing is that the soft soft interactions are not the driving forces of the reactions. They are
merely the consequences of the driving species that are in bonding process.
In the case of Lewis acids which are the transition metal ions that appear to expand their octets the
bonding may be actually stabilized by the soft soft interactions which lead to covalent bonding
accompanied bybonding as well .At present this is only a speculation !
QUESTIONS
1. Select the correct alternative for the following:
(i) Pearson’s Principle state……….

(a) Hard prefers to bind hard.
(b) Hard prefers to bind soft.
(c) soft prefers to bind hard.
(d) Hard prefers to bind hard soft prefers to soft.
(ii) Li+
is the….
(a) Soft base. (b) hard base
(c)soft acid (d) hard acid
(iii) Classification of acids and bases as HASB is due to…..
(a) Lewis (b) Arrhenius
(c) Bronsted (d) Pearson
(iv) For hard acids, KfΞ F-
>>Cl-
>Br-
>I-
(a) False (b) reverse is true
(c)For soft acids (d) correct
(v) H+
is called as hard acid whereas, H-
is a soft base.
(a) Faslse
(b) reverse is true
(c) H+
called soft acid
(d) correct

vi) Which of the following is hard acid
a) Pt2
b)Ag+
c) Hg2+
d) Li+
Which of the following is soft base a)NH3
b)H2O c)OH-
d)I-
vii) is the borderlined acid
a)Ag+
b)Al3+
C)k+
d)Zn2+
ix)----
Complex ion is stable
a)[AgF2]-
b)[Cd(NH3)4
2+
c)[CoI6]3-
d)[AgI2]-
[Ans: aleternate (d) for all].
2. Define the following :
(i) Pearson’s rule (ii) Hard acid
(iii) Soft base (iv) Boarderline acid v)Hard basevi)Soft acid
3. Write short note on :
(i) Pearson’s Principle
(ii) Hard/Soft base
(iii) Hard/Soft acid
(iv) Limitation of pearson’s rule
(v) Applications of HSAB concept
4. Give a brief account of classification of acid as hard and soft.
5. Give an account of acid-base strength and hardness-softness.
6. How do you classify bases as hard and soft according to pearson’s rule.
7. Write precisely on classification of acids and bases as hard and soft.

8. Comment on HSAB concept with suitable examples.
9. What do you know about acids-base strength and hardness- softness.
10. Write a critical note on applications and limitations of HSAB principle.
11. What do you mean by Hard and Soft acids and bases?Give their characteristics and
classification.
12. Discuss the therotical bases of hardness and softness.How are they related to acid-
base strength and electronegativity
13. According to pearson’s rule distinguish between a)Hard acid and soft acid b) Hard
base and soft base
.

MODULE-3 NON AQUEOUS SOLVENTS
1.3 Chemistry of Non aqueous Solvents.
1.3.1 Introduction, definition and characteristics of solvents.
1.3.2 Classification of solvents.
1.3.3 Physical properties and Acid-Base reactions in Liquid Ammonia (NH3) and
Liquid Sulphur Dioxide (SO2).
Introduction: An inorganic nonaqueous solvent is a solvent other than water that is not an organic compound.
These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special
environment.
Solution is a homogeneous mixture of two or more substances. A Solvent is a medium which dissolves a
solute to generate a solution .The dissolve substances in a solvent to produce solution is called a solute. The
substance that dissolves other substances to produce solution is called a solute.
Difference between Solute and Solvent
The basic difference between solute and solvent is that the former dissolves and the later is a dissolving
medium. The main difference between solute and solvent is that a solute is a substance that is added to a
solvent to form a solution. A solvent is a substance that dissolves the solute particles during the formation of a
solution. Let us now understand more about the difference between solute and solvent by studying in detail
What is Solute?
A solute is a substance that can be dissolved by a solvent to create a solution. A solute can come in many
forms. It can be gas, liquid, or solid. The solvent, or substance that dissolves the solute, breaks the solute apart
and distributes the solute molecules equally. This creates a homogenous mixture.
What is Solvent?
A material in which solute dissolves, resulting in a solution is a solvent. We always find solvent as a liquid but
it can also be a solid, a gas, or a supercritical fluid.
Parameter Solute Solvent

Meaning
A substance that gets
dissolved.
It is a dissolving medium.
Boiling point Higher than solvent Lower than solute
Dependability
Solubility depends on the
properties of the solute.
Solubility depends on the
properties of the solvent.
Physical state
Found in solid, liquid, or
gaseous state.
Found mainly in the liquid
state, but can be gaseous as
well.
Solute
We define solute as a substance that is dissolved in a solvent. A solute is a component of a solvent which, upon
getting dissolved, changes its form and loses its original characteristics. A solute is usually in smaller amounts
in the solvent. A common example of solute is salt and water. Salt dissolves in water and therefore, salt is the
solute.
Solvent
The solvent is a liquid, in which other materials dissolve to form a solution. Now Polar solvents like water,
favor the formation of ions while nonpolar ones like hydrocarbons do not. Solvents may be acidic, basic,
amphoteric (both), or aprotic (neither).
Classification of Solvents : Solvents can be classified is based on physical and chemical properties. These
are as:
i.) Liquid ,solid and gaseous solvents.
ii.) Protic and aprotic solvents.
iii.) Ionizing and non ionizing solvents.
iv.) Aqueous and non aqueous solvents.

i) Liquid ,solid and gaseous solvents: This classification is based on the physical state of solvents
.
ii) Protic and aprotic solvents.: It is based on chemical composition .
A. Protic or Protonic Solvents.: The solvent has at least a one hydrogen atoms in its solvents .
it is auto ionized or solvated proton.
eg. NH3 and H2O
NH3 + NH3 NH2
-
+ H+
( Solvated )
NH3 + H+
NH4
+
Protic or protonic solvents :- 2NH₃⇌ NH₄+ (ammonium) + NH₂− (amide)
3HF ⇌ H₂F+ + HF²- (hydrogen difluoride)
2H₂SO₄⇌ H₃SO₄+ + HSO₄- •
(A)Acidic or protogenic solvents. H₂SO₄,HCL,CH₃COOH,HCN. •
(B) Basic or protophilic solvents. NH₃, N₂H₄.
• Aprotic or non- protonicsolvents:- • C₆H₆,CHCl₃,CCl₄,SO₂.
N₂O₄ ⇌ NO+ (nitrosonium) + NO₃− (nitrate)
2SbCl₂⇌ SbCl₂+ (dichloroantimonium) + SbCl₄- (tetrachloroantimonate) POCl₃⇌ POCl₂+
+ POCl₄-
2 . acidic ,basic or amphiprotic solvents:- the solvents which have a tendency to donate
protons are acidic in nature and are called acidic solvents .eg.acetic acid ,hf. the solvents which
have strong tendency to accept protons are basic in nature and are called basic or protophilic
solvents.eg.liquid nh₃,c₆h₅n(pyridine) etc. the solvents which neither have tendency to gain nor
to lose protons are called amphiprotic or amphoteric solvents.eg.h₂o,c₂h₅oh, etc.
3.ionising and non ionising solvents:- • the solvents which are capable of undergoing self
ionisation(auto-ionisation)are called ionising solvents.eg.h₂o,nh₃,so₂, etc. • the solvents which
do not ionise at all • are called non-ionising solvents.eg. • benzene,hydrocarbons etc. • -these
are non-polar in nature.
4.co-ordinating and non co-ordinating solvents :- • the solvents which are capable of
coordinating with the metal ions or anions of the solute are called co-ordinating solvents. for
eg.nh₃,so₂,dmso, dmf etc. • on the other hand,the solvents which are not capable of co-

ordinating with the metal ions • of solute are called non co-ordinating solvents. • for eg.ccl₄,
saturated hydrocarbons etc.
question:- • which of the following are amphiprotic solvents? 1 h₂so₄ 2 hcl 3 h₂o 4 chcl₃
question:- • which of the following are example of protonic solvent? 1 hcn 2 chcl₃ 3 so₂ 4
ccl₄
mn(vii) cr(vi) v(v)
physical properties of solvents:-
• melting point and boiling point.
• dielectric constant.
• viscosity.
• dipole moment.
• heat of fusion & heat of vaporisation.
physical properties:-
• liquid ammonia as non-aqueous solvent:-
• freezing point:- -77.7 ċ
• boiling point:- -33.38 ċ
• dielectric constant:- 22.0 at -33.5 ċ
• liquid range:- -77 to -33 ċ
• heat of fusion:- 0.018 kj mol¯¹
• heat of vaporisation:- 23.6 kj mol¯¹
• self ionisable in nature • acts as an associated solvent
why ammonia acts as a better solvent
than water:- • poor conductor of electricity
• specific heat of ammonia is greater than
• water
• less viscous than water
• high critical temperature and pressure
• less associated than water(due to lesser
• hydrogen bonding)

• due to formation of strong reducing
• metal –ammonia solutions with aikali
• metals.
chemical reaction
• acid base reaction.
• precipitation reaction.
• redox reaction.
• solvation reaction.
type of reactions in non aqueous solvents:-
1) metathetical or precipitation reactions:- the reactions in which precipitation occurs on
mixing two solutions are called metathetical or precipitation reactions.
2AgNO₃ + BaCL₂ → 2AgCl ↓ +Ba(NO₃)₂
2) salt formation :-
• the reactions between appropriate acidic and basic substances to form salts are called salt
formation reactions.for eg,sodium ureide can not be prepared by the action of urea on sodium
hydroxide in water(because strong base can not take proton from urea molecule).
Na⁺(NHCONH₂)⁻ + H₂O → Na⁺ OH⁻ + NH₂CONH₂ (sodium ureide) (urea) however,this can
be easily formed in liq.nh₃ by reaction of urea with sodamide. NH₂CONH₂ + Na⁺NH₂⁻→
Na⁺(NHCONH₂ )⁻ +NH₃ (urea)(sodamide) (sodium ureide)
3)acid base reactions :-
• acid base reaction can be explained on the basis of solvent system concept,
• an acid is a substance that by direct dissociation or reaction with the solvent gives the cation
chacteristic of the solvent.similarly,
• a base is a substance that gives the anion characteristic of the solvent.
• for eg, in liq. NH₃ solvent ,NH₄⁺ ion act as acid and NH₂⁻ ions act as base. the neutralisation
reaction is:- • NH₄Cl + NaNH₂ → NaCl + 2NH₃ acid base salt solvent
3) solvolytic reactions or solvolysis :-

• the reactions in which the solvent molecules react with the solute in such a way that the
solvent molecules split up into two parts, one or both of which get attached to the solute
molecule or ion are called solvolytic reactions.
• water as solvent:- 2H₂O↔ H₃O⁺+ OH⁻ (AUTO IONISATION )
• SO₂CL₂ + 4 H₂O↔ SO₂ (OH)₂ + 2H₃O⁺ +2CL⁻
• F⁻ +H₂O ↔ HF +OH⁻
• AMMONIA AS SOLVENT :- 2NH₃↔ NH₄⁺ + NH₂⁻
• SO₂CL₂ + 4 NH₃ ↔ SO₂(NH₂)₂ + 2NH₄⁺ + 2CL⁻
• H⁻ + NH₃↔ NH₂⁻ + H₂
4) solvation reaction :-
5) • solvent get attached to a solute species( cation , anion , or molecule) are called solvation
reactions.the species formed is called solvate.
6) • CuSO₄ + 4NH₃ → CuSO₄.4NH₃ (ammoniate) • CuCl₂ + 4 H₂O → [ Cu(H₂O)₄]²⁺ +2 Cl⁻
(HYDRATE) • BaSO₄ + 3 H₂ SO₄→ BaSO₄.3H₂SO₄ (solvate of sulphuric acid)
reactions in liquid ammonia 1)acid base reactions:-
liquid ammonia ionises as:- 2NH₃↔NH₄⁺ + NH₂⁻
for eg.
KNH₂→ K⁺ +NH₂⁻ protolysis reaction:- certain compounds like urea which are incapable of
donating pprotons to water can readily donate proton to ammonia in liquid amonia . NH₂ NH⁻
C = O + NH₃ → C = O + NH₄⁺ ⁄ ⁄ NH₂ NH₂
2) acid base neutralisation reaction:- • neutralisation of an acid and base in liquid ammonia
involves combination of NH₄⁺ ion(from acid) and nh₂⁻ ion(from base):- NH₄CI + KNH₂→
KCI +2NH₃ ACID BASE NH₄⁺ + NH₂⁻ → 2NH₃
• AS AN ACID:- ZN(OH)₂ + KOH → K₂ZN(NH₂)₄
• AS A BASE:- ZN(NH₂)₂ +2NH₄CI→ ZNCI₂ + 4NH₃
3) precipitation reactions:- • precipitation reactions involve double decomposition because
of the differences in solubilities.
KCl +AgNO₃→ AgCl + KNO₃ • white ppt of BaCl₂ is produced when silver chloride and
liq. ammonia brought together :- 2AgCl+ Ba(NO₃)₂↔ BaCl₂ (PPT)+ 2AgNO₃

4) ammonolysis reaction:- • the solvolysis reactions in liquid ammonia are called
ammonolysis or ammonolytic reactions. • for eg.hydrolysis of SiCl₄ occurs as:-
• SiCl₄ +4H₂O ↔ Si(OH)₄ (silicic acid) • organic halides undergo slow ammonolysis reaction
to form ammines :- RX + 2NH₃ → RNH₂ + NH₄X (primary amines)
5) solutions in liquid ammonia:-
• the most striking property of liquid ammmonia is its ability to dissolve alkali metals.the
resulting solutions are blue and good electrical conductors. • when alkali metals are dissolved
in liquid ammonia they ionise to give metal ions and valence electrons as:- Na → Na⁺ + E⁻ •
both alkali metal and electron become solvated by ammonia molecules.
NA⁺ + XNH₃→ [ NA(NH₃)X]⁺ (ammoniated cation) • E⁻ + YNH₃ → [ E(NH₃)Y]⁻
(amnoniated electron ) • the complete reaction may be written as:- na → [ na⁺ (nh₃)x]⁺ + [ e
(nh₃)y]⁻ • the ammoniated electrons are responsible for blue colour of solution.
liquid sulphur dioxide :- • it is a non –protonic solvent or aprotic solvent because it does not
contain any hydrogen atom. • it is also one of the important non aqueous solvent and widely
used in industry. • physical properties of liquid so₂ :- • freezing point :- -75.46 ċ • boiling poin
t :- -10.02 ċ • dielectric constant:- 17.40
reactions in liquid so₂ :-
1) acid base raction :- liquid SO₂ undergo auto ionisation as :
2SO₂ ↔ SO²⁺ + SO₃²⁻
acid base or neutralisation reaction in liquid SO₂
May be given as:- SOCl₂ + Cs₂SO₃ → 2CsCl + 2SO₂ (ACID) (BASE)
2) solvolytic reactions :- • solvolytic reactions in liquid so₂ are not so common as they are in
other solvents such as nh₃. • certain covalent halides such as pcl₅ ,pbr₅, wcl₆ under go
solvolysis in liquid so₂ in sealed tubes. • PCL₅ + SO₂ (LIQ) → POCL₃ + SOCL₂ • PBR₅ +
SO₂(LIQ) → POBR₃ + SOBR₂ • WCL₆ +SO₂ (LIQ) → WOCL₄ + SOCL₂
3) precipitation reactions :- • precipitation of several insoluble compounds in liq. sulphur
dioxide can be carried out by treating with soluble compound in so₂. sbcl₃ + 3lii → sbi₃ ↓ +
3LICL ALCL₃ + 3NAI → 3NACL ↓ + ALI₃

4) redox reactions :- • liq. sulphur dioxide does show any marked reducing or oxidising
property. it simply act as a medium for certain redox reactions. 6ki + 3sbcl₅ → 2k₃[sbcl₆] +
sbcl₃ + 3i₂
Liquid Sulphur Dioxide
Liquid sulphur dioxide is also a non-protonic solvent as it does not yield a proton (H+
) on ionization. It is widely
used for carrying out a number of chemical reactions Under normal temperature and pressure, sulphur dioxide is a
gas but it can be readily liquefied. It has a wide liquid range (- 10o
C to -75.5o
C) and hence can serve as a good
solvent. Its use as a non-aqueous solvent is rapidly increasing due to its low cost and ease of its handling. Its
dielectric constant is small (17.4 at -20.0o
C) which makes it a good solvent for electrovalent compounds.The
characteristic physical properties of liquid sulphur dioxide are given in Table.Table: Physical Properties of Liquid
Sulphur Dioxide.
Properties Values
Boiling point -10.1o
C
Freezing point -75.5o
C
Density 1.46 g ml-1
(-10o
C)
Dielectric constant 17.4 (-20o
C)
Specific conductance
(ohm-1
cm-1
)
4x10-8
(-10o
C)
Viscosity (centipoise) 0.428 (-10o
C)
Dipole moment
(Debye)
1.61
Autoionisation.
By analogy with water and liquid ammonia, autoionisation of sulphur dioxide takes place as under:
SO2 + SO2 SO2+
+ SO3
2-
The thionyl ion (SO2+
) is analogous to the hydronium ion (H3O+
) and ammonium ion (NH4
+
) while sulphite ion
(SO3
2-
) corresponds to hydroxyl ion (OH-
) and amide ion NH2
-
of the aqueous and liquid ammonia systems,
respectively.
Sulphur dioxide solutions are not as good electrical conductors as are liquid ammonia or aqueous solutions.

Solubility of Substances in Liquid Sulphur Dioxide.
Amongst the inorganic compounds, iodides and thiocyanates
are the most soluble. Metal sulphates, sulphides, oxides and hydroxides are practically insoluble.
Many of the ammonium, thallium and mercuric salts are soluble. Liquid sulphur dioxide is an excellent solvent for
covalent compounds. Substances such as IBr, PBr3, CCl4, SiCl4, SnCl4 are soluble in it. Metals are insoluble in
liquid sulphur dioxide.Amongst the organic compounds, benzene and alkenes dissolve in it freely. Pyridine,
quinoline, ethers, halogen derivatives and acid chlorides also dissolve in liquid sulphur dioxide. Alkanes are
insoluble.Reactions in Liquid Sulphur Dioxide. Chemical reactions that occur in liquid sulphur dioxide are of the
following types:
1. Acid-base Reactions or Neutralisation Reactions. Comparing the autoionisation of liquid sulphur dioxide
with that of water, it can be seen that thionyl ion (SO2+
) is analogous to hydronium ion and sulphite ion (SO3
2-
) is analogous to hydroxyl ion (OH-
). Hence, all compounds containing or making available SO3
2-
ions in
liquid sulphur dioxide will act as bases in this medium. Similarly, all compounds which make
available SO2+
ions in liquid sulphur dioxide will act as typical acids in liquid sulphur dioxide.
Typical acid-base or neutralisation reactions in liquid sulphur dioxide are given below.
Reaction between thionyl chloride and cesium sulphite:
liq. SO2
SOCl2 + Cs2SO3 2CsCl + 2SO2
Acid Base Salt Solvent
Reaction between thionyl thiocynate and potassium sulphite.
liq. SO2
SO(SCN)2 + K2SO3 2K(SCN) + 2SO2
Acid Base Salt Solvent
2. Solvolytic Reactions. Only a limited number of salts undergo solvolysis in liquid sulphur dioxide. Some
common reactions are given below.
Ammonium acetate is solvolysed in liquid sulphur dioxide.
2CH3COONH4 + 2SO2 (NH4)2SO3 + (CH3COO)2SO (CH3COO)2SO

SO2 + (CH3CO)2O
Binary halides such as PCl5, UCl6, WCl6 undergo solvolysis in liquid sulphur dioxide.
PCl5 + SO2 POCl3 + SOCl2 UCl6
+ 2SO2 UO2Cl2 + 2SOCl2
WCl6 + SO2 WOCl4 + SOCl2
The formation of solvates, i.e., the addition compounds with the solvent, is also known. Typical solvates formed are
NaI.4SO2, RbI.4SO2, KBr.4SO2, CaI2.4SO2, BaI2.4SO2, SrI2.4SO2, AlCl3.2SO2, etc.
3. Precipitation Reactions. A large number of precipitation reactions can be carried out in liquid sulphur
dioxide due to specific solubility relationships. Some of these reactions are given below.
liq. SO2
2CH3COOAg + SOCl2 2AgCl + SO(CH3COO)2
Thionyl acetate
liq. SO2
2KI + SOCl2 2KCl + SOI2
liq. SO2
SbCl3 + 3LiI SbI3 + 3LiCl
liq. SO2
PbF2 + Li2SO4 PbSO4 + 2LiF
liq. SO2
AlCl3 + 3NaI 3NaCl + AlI3
4. Complex Formation Reactions. A large number of complex formation reactions of liquid sulphur dioxide
have been reported. For instance, the solubility of iodine in liquid sulphur dioxide is greatly increased by the
addition of potassium or rubidium iodide. This is due to the formation of the complex KI3 or RbI3.

liq. SO2
KI + I2 KI3
liq. SO2
RbI + I2 RbI3
Similarly, the increase in the solubility of cadmium iodide and mercuric iodide in liquid sulphur dioxide is
attributed to the formation of complexes.
liq. SO2
HgI2 + 2KI K2[HgI4]
5. Amphoteric Behaviour. Various salts show amphoteric behaviour in liquid sulphur dioxide. The reaction
of AlCl3 with NaOH in aqueous medium can be compared with the reaction of AlCl3 with tetramethyl
ammonium sulphite in liquid sulphur dioxide. In aqueous solution, the reaction takes place as follows.
H2O
AlCl3 + 3NaOH Al(OH)3 + 3NaCl
Gelatinous ppt.
H2O
Al(OH)3 + NaOH Na[Al(OH)4]
Soluble complex
From the soluble complex, Al(OH)3 can be reprecipitated by the addition of HCl.
Na[Al(OH)4] + HCl NaCl + H2O + Al(OH)3
In liquid sulphur dioxide medium, an identical reaction takes place between AlCl3 and tetramethyl ammonium
sulphite:
liq. SO2
2AlCl3 + 3[N(CH3)4]2SO3 6[N(CH3)4]Cl + Al2(SO)3
Gelatinous ppt.
liq.SO2

Al2(SO3)3 + 3[N(CH3)4]2SO3 2[N(CH3)4]3[Al(SO3)3]
Soluble Complex From the
soluble complex, Al2(SO3)3 can be reprecipitated by adding
the acid, SOCl2.
liq.SO2
2[N(CH3)4]3[Al(SO3)3] + 3SOCl2 6[N(CH3)4]Cl + 6SO2 + Al2(SO3)3
The behaviour of GaCl3 is similar to that of AlCl3.
6. Redox Reactions. Liquid sulphur dioxide does not have any strong oxidising or reducing properties. It serves
only as a medium for redox reactions. For instance, liquid sulphur dioxide cannot reduce iodine. However, a
sulphite in liquid sulphur dioxide reduces iodine to iodide.
liq. SO2
I2 + 2R3SO3 R2SO4 + 2RI + SO2
KI is oxidised to free iodine by SbCl5 in liquid sulphur dioxide.
liq. SO2
6KI + 3SbCl5 3I2 + SbCl3 + 2K3[SbCl6]
Thank You

Module -1 -3 Acid bases and non aqueous solvents

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Module -1 -3 Acid bases and non aqueous solvents