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According to Lewis Concept an Acid-Base reaction involves an interaction of a vacant orbital of an acid (A) and a filled or unshared orbital of a base (B).
A + :B A : B 
Lewis Acid

Lewis Base

Adduct or Complex
The species A is called Lewis Acid or a generalized acid and B is called Lewis Base or a generalized base. A strong acid A and a strong base B, will form the stable complex A : B.
A concept known as Principle of Soft and Hard Acids and Bases(SHAB) Principle is very helpful in making a stability of the complex A : B.
According to this principle the complex A : B is most stable when A and B are either both soft or both hard.
The complex is least stable when one of the reactants (namely A and B) is very hard and the other one is very soft. 
In order to arrive at a comparative estimate of the donor properties of different bases, the preferences of a particular base to bind a proton H+ and methyl mercury (II) ion, [CH3HgB]+ was determined. Both the proton and methyl mercury cation can accommodate only one coordinate bond but the two cation vary widely in their preferences to bases. This preferences was estimated from the experimental determination of equilibrium constants for the exchange reactions:
BH+ + [CH3Hg(H2O)]+ [CH3HgB]+  + H3O+
The results indicate that bases in which the donor atom is N, O or F prefer to coordinate to the proton. Bases in which the donor atoms is P, S, I, Br, Cl or C prefer to coordinate to mercury.

Hard and Soft Bases:

The donor atoms in the first group have high Electronegativity, low Polaris ability and hard to oxidize. Such donors have been named ‘hard bases by Pearson, since they hold on to their electrons strongly. 
The donor atoms of second category are of low electronegativity, high polarisability, and are easy to oxidize. Such donors have been called ‘soft’ bases since they are hold on to their valence electrons rather loosely.
In simple terms hardness is associated with a tightly held electron shell with little tendency to polarise. On the other hand softness is associated with a loosely bound polarisable electron shell.
It will be seen that within a group of the periodic table softness of the Lewis bases increases with the increase in size of the donor atoms. 
Thus, among the halide ions softness increases in the order:
-Cl -Br -I -
Thus F - is the hardest and I - is the Softest base.
Definition of hard acid and soft acids and bases
Classification of Bases

Hard and Soft Acids:

After having gone through a classification of bases, a classification of Lewis acids is necessary. The preferences of a given Lewis acid towards ligands of different donor atoms is usually determined from the stability constant values of the respective complexes or from some other useful equilibrium constant measurements. When this is done, metal complexes with different donor atoms can be classified into two sets based on the sequences of their stabilities.
Hard acids have small acceptor atoms, are of high positive charge and do not contain unshared pair of electrons in their valence shell, although all these properties may not appear in one and the same acid. These properties lead to high electronegativity and low polarisability. In keeping with the naming of the bases, such acids are termed as 'Hard'Acids.
Hard Acids:

Soft acids have large acceptor atoms, are of low positive charge and contain ushered pairs of electrons in their valence shell. These properties lead to high polarisability and low electronegativity. Again in keeping with the naming of the bases, such acids are termed 'Soft' Acids
Soft Acids:
Difference between hard acid and soft acid and bases
Classification of Lewis Acids as Hard, Intermediate and Soft Acids
Classify the following as Hard and Soft Acids and Bases. 
(i) H- (ii) Ni+4 (iii) I+ (iv) H+
  1. The hydride ion has a negative charge and is far too large in size compared to the hydrogen atom. Its electronegativity is quite low and it will be highly polarisable by virtue of its large size. Hence it is Soft Base.
  2. Quadrivalent nickel has quite a high positive charge. Compared to bivalent nickel its size will be much smaller. Its electronegativity will be very high and polarisability will be low. Hence it is Hard Acid.
  3. Mono-positive iodine has a low positive charge and has a large size. It has a low electronegativity and a high polarisability. Hence it is a Soft Acids.
  4. H+ has the smallest size with a high positive charge density. It has no unshared pair of electrons in its valence shell. All these will give a high electronegativity and very low polarisability. Hence H+ is a Hard Acid.
Soft and Hard Acids and Bases (SHAB) Principle:
This principle also means that if there is a choice of reaction between an Acid and two Bases and two Acids and a Base,
A Hard Acid will prefer to combine with a Hard Base and a Soft Acid will prefer to combine with Soft Base and thus a more stable product will be obtained.
Hard Acid - Hard Base may interact by strong ionic forces. Hard Acids have small acceptor atoms and positive charge while the Hard Bases have small donor atoms but often with negative charge. Hence a strong ionic interaction will lead to Hard Acid - Base combination.
On other hand a Soft Acid - Soft Base combination mainly a covalent interaction. Soft Acids have large acceptor atoms, are of low positive charge and contain ushered pair of electrons in their valence shell.

Application of Soft and Hard Acids and Bases (SHAB) Principle:

  1. [CoF6]-3 is more stable than [CoI6]-3
  2. It will be seen that Co+3 is a hard acid
    F- is a hard base and I- is a soft base
    Hence [CoF6]-3 (Hard Acid + Hard Base) is more stable than [CoI6]-3 (Hard Acid + Soft Base).
  3. The existence of certain metal ores can also be rationalised by applying SHAB principle. Thus hard acids such as Mg+2, Ca+2 and Al+3 occur in nature as MgCO3, CaCO3 and Al2O3 and not as sulphides (MgS, CaS and Al2S3), since the anion CO3-2 and O-2 are hard bases and S-2 is a soft base.
  4. Soft acids such as Cu+, Ag+ and Hg+2 on the other hand occurs in nature as sulphides. The borderline acids such as Ni+2, Cu+2 and Pb+2 occur in nature both as carbonates and sulphides. 
    The combination of hard acids and hard bases occurs mainly through ionic bonding as in Mg(OH)2 and that of soft acids and soft bases occurs mainly by covalent bonding as in HgI2.
AgI2- is stable, but AgF2- does not exist. Explain.
We know that Ag+ is a soft acid, F- is hard base and I- is soft base.
Hence AgI2- (Soft Acid + Soft Base) is a stable complex and AgF2- (Soft Acid + Hard Base) does not exist.
Explains why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not?
In the case of Hg(OH)2 and HgS, Hg is a soft acid and OH- and S-2 is hard base and soft base respectively. Evidently HgS (Soft acid + Soft base) will be more stable than Hg(OH)2 (Soft Acid + Hard base). More stability of HgS than that of Hg(OH)₂ explain why Hg(OH)2 dissolved readily in Acidic Aqueous Solution but HgS does not. 


Any of the class of substances whose aqueous solutions are characterized by the sour taste, the ability to turns blue litmus red, and the ability to react with bases and certain metals to form salts.


In chemistry, a base is a substance that in aqueous solution is slippery to touch, that's bitter, changed the color of the indicator, (turns red litmus to blue) react with acid to form salts.
The concepts of acid and base is developed one after another tend to make the definitions more and more broad based.
We need to familiar with the following five concepts :
1. The Arrhenius Concept.
2. The Solvent System Concept.
3. The Protonic Concept.
4. The electronic Theory.
1. The Arrhenius Concept :
Arrhenius was one of the early exponent of electrolytic dissociation theory to define acids and bases.
His classification of acids and bases was based on the theory that, acids when dissolved in water dissociate hydrogen ions and anions where as bases when dissolved in water dissociate into hydroxyl ions and cations. Thus HCl is an acid and NaOH is a base and the neutralization process can be represented by a reaction involving the combination of H⁺ and OH⁻ ions to form H₂O.
Acid :   HCl H⁺ + Cl⁻ 
Base :   NaOH Na⁺ + OH⁻
Acid Base Equilibrium: H⁺ + OH⁻  H₂O
Acid base neutralization reaction in aqueous medium :
The Arrhenius concept was highly useful to explaining acid base reaction in aqueous medium.
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Representation of Acid Base Neutralization Reaction
The heat liberated during neutralization of strong acid (HCl, HClO₄, HNO₃, HBr, HI and H₂SO₄) and strong base (NaOH, KOH, RbOH, and Ca(OH)₂) is the same, namely 13.4 kcal/mole (56 KJ/mole) indicating that the ion participating in the reaction must be the same (that is H⁺ and OH⁻) and that other ions (Na⁺, K⁺, Rb⁺, Ca⁺², Cl⁻, Br⁻, I⁻, ClO₄⁻ or SO₄⁻²) take no part.
This can be explain as follows.Let, 
HA(aq)+BOH(aq) BA(aq)+H₂O
Where, HA and BOH are strongly electrolytes and completely dissociate in the Solution.
Hence, (H⁺+A⁻)+( B⁺+OH⁻)(B⁺+A⁻)+ H₂O
Cancelling likes ions, we have,
H⁺+OH⁻  H₂O
Since neutralization of all strong acids and bases reduced to the formation of 1 mole water from H⁺ ion and OH⁻ ion.The enthalpy change (ΔH) will also be same.
Limitation of Arrhenius Concept:
(i) According to this concept, HCl is regarded as an acid only when dissolved in water and not in some other solvent such as benzene or when it exists in the gaseous state.
(ii) It cannot account for the acidic and basic character of the materials in non aqueous solvents, as for example, NH₄NO₃, in liq NH₃ acts as an acid, through it does not give H⁺ ions. Similarly many organic materials in NH₃ , which does not give OH⁻ ions at all, are actually known to show basic character.
(iii) The neutralization process limited to those reactions which can occurs in aqueous solutions only, although the reactions involving salt formation do occur in many other solvents and even in the absence of solvents.
(iv) It cannot explain the acidic character of certain salts such as AlCl₃ in aqueous solution.
2. The Solvent system Concept:
The non- aqueous solvent molecules may also dissociate into two oppositely charged ions.
(i)We consider the solvent H₂O, its characteristic cation and anion are H⁺ and OH⁻ respectively are, 
H₂O H⁺ + OH⁻
Since we know that a bare proton will readily polarize other anions or molecules we write a H⁺ as H₃O⁺ indicating that it is a solvated proton that exists in the solution.
So that the overall dissociation of solvent will be:
H₂O + H₂O H₃O⁺ + OH⁻
Thus all those compounds which can give H₃O⁺ ions in H₂O will act as an acids and all those compounds which can give OH⁻ ions in H₂O will behave as bases.
(ii) Similarly liquid NH₃ can written as 
NH₃ + NH₃ NH₄⁺ + NH₂⁻
Thus those compounds which give NH₄⁺ ions in liquid NH₃ will act as an acids and all those compounds which can give NH₂⁻ ions in liquid NH₃ acts as bases.
Thus the dissociation (or autoionisation) of non-aqueous solvents is directly responsible for the nature of the chemical reactions that can be initiated in such solvents.
According to solvent system concept,
An acid is a substance which by dissociation in the solvent forms the same cation as does the solvent itself due to auto-ionization.
A base is one that, gives on dissociation in the solvent the same anion as does the solvent itself on its ionization.
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Acid-Base Neutralisation Reaction According to Solvent System Concept
The auto-ionisation of some protonic and non-protonic solvents are, Protonic Solvents:
H₂O + H₂O H₃O⁺ + OH⁻
NH₃ + NH₃ NH₄⁺ + NH₂⁻
Non-Protonic Solvents:
SO₂ + SO₂ SO⁺² + SO₃⁻²
BrF₃ + BrF₃ BrF₂⁺ + BrF₄⁻
N₂O₄ + N₂O₄ 2NO⁺ + 2NO₃⁻
Neutralization Reactions according to Solvent System Concept: 
Just as with the Arrhenius definition, neutralisation is a reaction between an acid and a base to produce salt and solvent. 
Neutralisation of some non-aqueous solvents are, 
(i) Neutralisation reaction in liquid NH₃: 
Dissociation of Solvent:    
NH₃ + NH₃  NH₄⁺ + NH₂⁻
Dissociation of Acid:         
NH₄Cl NH₄⁺ + Cl⁻ 
Dissociation of Base:         
KNH₂  K⁺ + NH₂⁻
Acid Base Neutralization reaction: 
NH₄Cl + KNH₂ KCl + 2NH₃
NH₄Cl(Acid) + KNH₂(Base)  KCl(Salt) + 2NH₃(Solvent)
(ii) Neutralization reaction in liquid SO₂:
Dissociation of Solvent:
SO₂SO₂  SO⁺² + SO₃⁻²
Dissociation of Acid:
SOBr₂ SO⁺² + 2Br⁻ 
Dissociation of Base: 
K₂SO₃ 2K⁺ + SO₃⁻²
Acid Base Neutralization reaction:
SOBr₂ + K₂SO₃ 2KBr + 2SO₂
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Neutralization Reactions according to Solvent System Concept

Utility of the Concept:

Evidently this concept of solvent system can be used to explain the acid-base reactions occurring in aqueous and non aqueous solvents(protonic or non-protonic both).
Thus this theory can simply be said to be an extension of the Arrhenius Theory.
(i) This theory does not consider a number of acid-base reactions included in the protonic definition.
(ii) It limits acid-base phenomena to solvent system only. Thus it does not explain the acid base reactions which may occurs in the absence of solvent.
(iii) It can not explain the neutralisation reactions occurring without the presence of ions.

Each of the concepts had its own limitations. The need was therefore felt for a concept which can cover all the earlier aspects of Acids and Bases. Lewis, (1923) explains the acid-base phenomenon not in terms of ionic reactions but in terms of electronic structure of the acid and base along with the formation of the coordinate covalent bond.
A base was to be considered according to Lewis as an electron pair donor and an acid as an electron pair acceptor.
An acid therefore combines with a base leading to a coordinate link from the base to the acid.
An acid must have a vacant orbital into which an electron pair donated by a base can be accommodated. A base, then, is a substance which has at least one unshared (lone) pair of electrons.
Vacant 1S
Lone Pair of
H+ : NH3
Acid Base
According to Lewis theory, the process of neutralisation is simply the formation of a coordinate bond between an acid and a base. 
The neutralisation product, termed as coordinate complex or adduct, may be either non-ionisable or may undergoes dissociation or condensation reaction depending on its stability.
H+ + :NH3 [H  :NH3]+
F3B + :NH3 F3B :NH3
Bronsted Concept Vs Lewis Concept:
According to Bronsted Concept, a base is any molecules or ion, which accepts proton. 
Lewis Concept recognises that the base is an electron pair to donate, H+
So that during neutralization a coordinate bond is formed between the donor atom of the base and the proton.
H+ + :NH3 [H  :NH3]+
Thus all Bronsted bases are bases in Lewis sense also.
Arrhenius Concept Vs Lewis Concept:
According to Arrhenius, an acid when dissolved in water dissociate into hydrogen ions and anions, that is it release hydrogen ions (H+). Since the proton can receive electron pair from a base. Thus all the Arrhenius Acids are also acids under Lewis definition.
Solvent System Concept Vs Lewis Concept:
The solvent system concept recognizes that ammonium chloride should behave as an acid in liquid ammonia. It may be noted that ammonium ion is but an ammoniated proton and the proton can accept a lone pair of electron from a base. Therefore an ammonium ion in liquid ammonia is also an acid in Lewis sense.
The solvent system considers amide ion is a base in liquid ammonia.The amide ion is capable of donating a lone pair to an acid, hence this is a base in Lewis sense.
The reaction between the ammonium ion and amide ion in liquid ammonia is represented in Lewis sense as follows:
NH3 + [H+ NH2-]  NH4+  + NH2-

Classification of Lewis Acids and Bases: 

A wide variety of compounds are recognized as acids and bases according to Lewis. Any Lewis acid must contain at least one empty orbital in the valence shell of one of its atoms to accept an electron pair from a Lewis base.
(i) Molecules containing a central atom with an incomplete octet:
Typical examples of this class of acids are electron deficient molecules such as alkyl and halides of Be, B and Al
Some reactions of this types of Lewis acid with Lewis base are,
F3B + O(C2H5)2 F3BO(C2H5)2

Cl3Al + NC5H5 Cl3AlNC5H5

Me3B + N2H4 Me3BN2H4

Explain - NH3 behaves as a base but BF3 as an Acid.
In NH3, the central N atom have lone pair of electrons This lone pair coordinate to empty orbital, is termed as a base according to Lewis concept. 
The compounds with less than an octet for the central atom are Lewis acids. In BF3B in BF3, contain six electrons in the central atom thus it behaves as an acid.
F3B + :NH3 F3BNH3
Explain why tri-covalent phosphorus compounds can serve both as Lewis acids and also as bases?
Tri-covalent phosphorus compounds like PCl3 have a lone pair of electrons in phosphorus. This lone pair may coordinate to a metal ion thus allowing the compound to serve as a Lewis base. 
Ni + 4PCl3 [Ni(PCl3)4]0
Again the quantum shell of phosphorus has provision for d-orbitals which can receive back donated electrons from electron reach low oxidation state of a metal ion. In this latter case tri-covalent phosphorus compound serves as a Lewis acid.
(ii) Molecules Containing a Central atom with Vacant d-Orbitals:
The central atom of the halides such as SiX4, GeX4, TiCl4, SnX4, PX3, PF5, SF4, SeF4, TeCl4, etc. have vacant d-orbitals. These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-orbital and can thus form adducts with a number of halide ions and organic bases. These substances is thus Lewis acids.
These halides are vigorously hydrolyzed the appropriate HX. The hydrolysis reaction take place through the formation of an unstable adduct(intermediate) with H2O.
Bi-positive tin can function both as a Lewis acid and a Lewis base. Explain?
The Lewis representation of SnCl2 shows a lone pair on tin through it is as yet short of an octet. Ligands, particularly donor solvents with lone pairs, may coordinate to tin giving complexes. Here SnCl2 behaves as a Lewis acid.
Again interaction of platinum group metal compounds with SnCl2 (SnCl3-) as donor leads to the formation of coordinate complexes. As for examples,
[(Ph3P)2PtCl(SnCl3)], [RuCl2(SnCl3)2]-2
These are examples of Lewis base behavior of SnCl2.
Can SiCl4 and SnCl4 function as Lewis acids?
Both silicon and tin are members of Group IV of the periodic table and their quantum shells admit of d-orbitals. As a result they can expand their valence shell through SP3d2 hybridisation and can give rise to six coordinate complexes.
In fact complex like [SiCl4{N(CH3)3}2], [SiCl2(Py)4]Cl2, [SnCl4(Bpy)] are known. Thus SiCl4 and SnCl4 can function as Lewis acids.
(iii) Simple Cations:
Theoretically all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis base are,
Lewis acid + Lewis base Adduct 

Ag+ + 2(:NH3) [NH3AgNH3]+

Cu+2 + 2(:NH3) [Cu(NH3)4]+2 

Co⁺³ + 6(:OH₂) [Co(OH₂)₆]⁺³ 

Li+ + :OHCH3 [LiOHCH3] 
The Lewis acid strength or coordinating ability of the simple cations are increases with- 
(a) An increases in the positive charge carried by the cation.
(b) An increases in the nuclear charge for atoms in any period of the periodic table.
(c) A decreases in ionic radius.
(d) A decreases in the number of the shielding shells. 
Evidently the acid strength of simple cations increases for the elements on moving from left to right in a period and from bottom to top in a group on the periodic table.
Strength of Lewis acid increasing:
( Positive charge increases from +2 to +3)
(On moving from bottom to top in a group)
(On moving from left to right in a period)
Arrange these oxides in order of their acidic nature : N2O5, As2O3, Na2O, MgO.
Acidic oxides react with water to give oxoacids. The higher the oxidation number and the higher the electronegativity the grater the central element will force to reaction with water to give the oxoacid. Of all these oxoacids nitrogen has the highest oxidation number and electronegativity. Thus the order of acidic nature and oxidation number are-
Acidic nature: 
Oxidation number of Oxide forming element: +1+2+3+5.
(iv) Molecules having a multiple bond between atoms of dissimilar electronegativity: 
Example of these class are CO2, SO2, and SO3. In these compounds the oxygen atom are more electronegative than sulphur or carbon atom. As a result the electron density of π-electrons is displaced away from carbon to sulphur atoms which are less electronegative than oxygen, towards the O-atom.Thus carbon and sulphur are electron deficient able to accept the electron pair from a Lewis base such as OH- ions to from dative bond.
What is the Lewis Concept of Acids and Bases?
Example of Lewis Acid Base Reaction
Bisulphate ion can be viewed both as an acid and a base. Explain.
Bisulphate ion is HSO3-. It may lose a proton to give the conjugate base SO3-2, thus behaving as an acid. Again it may add on a proton to give the conjugate acid, thus showing its base character.
SO3 Behaves as an acid and H2O as a base. Explain.
SO3 like BF3 has less than an octet,and will be termed as an acid according to Lewis concept.
But in H2O oxygen atom contain lone pairs to donate the Lewis acid.
(v) Elements with an Electron sextet:
Oxygen and sulphur contain six electrons in their valence shell and therefore regarded as Lewis Acids.
The oxidation of SO3-2 to SO4-2 ion by oxygen and S2O3-2 ion by sulphur are the acid-base reactions.
SO3-2 + [O] [OSO3]-2 
SO3-2 + [S] [SSO3]-2

Utility of Lewis Concept:

(i) This concept includes those reactions in which no proton are involved.
(ii) Lewis concept is more general than the Bronsted - Lowry Concept (that is Protonic Concept) in that acid base behavior is not dependent on the presence of one particular element or on the presence or absence of a solvent.
(iii) It explains the long accepted basic properties of metallic oxides and acidic properties of non-metallic oxides.
(iv) This theory explain many reactions such as gas-phase high temperature and non-solvent reaction as neutralisation process.

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