Lewis concept of acids and bases

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 the 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.
    H⁺: NH3
    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 undergo dissociation or condensation reaction depending on its stability.
H+ + : NH3 [H ← :NH₃]+
Acid Base Adduct
    F₃B + :NH₃ F₃B ← :NH₃

Bronsted Concept Vs Lewis Concept:

    According to Bronsted Concept, a base is any molecules or ion, which accepts a 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 releases hydrogen ions (H+). Since the proton can receive electron pair from a base. Thus all the Arrhenius Acids are also acids under the 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 an electron from a base. Therefore an ammonium ion in liquid ammonia is also 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.

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 these types of Lewis acid with Lewis base are,
F3B + O(C2H5)2 F3B←O(C2H5)2
Acid Base Adduct
Cl₃Al  + NC₅H₅   Cl₃Al ← NC₅H₅
Me₃B + N₂H₄  Me₃B ← N₂H₄

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 are thus Lewis acids.
    These halides are vigorously hydrolyzed the appropriate HX. The hydrolysis reaction takes place through the formation of an unstable adduct(intermediate) with H₂O.

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) ⇄ [NH₃→Ag←NH₃]⁺
Cu⁺² + 2(:NH3) ⇄ [Cu(NH3)4]⁺²
Co⁺³ + 6(: OH₂) ⇄ [Co(OH₂)₆]⁺³
Li⁺ +:  OHCH₃ ⇄ [Li←OHCH₃]
    The Lewis acid strength or coordinating ability of the simple cations are increased with-
  1. An increase in the positive charge carried by the cation.
  2. An increase in the nuclear charge for atoms in any period of the periodic table.
  3. A decrease in ionic radius.
  4. A decrease 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:
    Fe+2Fe+3 ( Positive charge increases from +2 to +3) K+Na+ (On moving from bottom to top in a group) Li+Be+2  (On moving from left to right in a period)

Molecules having multiple bonds between atoms of dissimilar electronegativity:

    Example of this class is CO₂, SO₂, and SO₃. In these compounds, the oxygen atom is 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 form a dative bond.
What is the Lewis Concept of Acids and Bases?
Example of Lewis Acid-Base Reaction

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] [O←SO₃]⁻²
SO₃⁻² + [S] [S←SO₃]⁻²

The utility of Lewis Concept:

  1. This concept includes those reactions in which no proton are involved.
  2. Lewis concept is more general than the Bronsted - Lowry Concept (that is Protonic Concept) in that acid-base behaviour is not dependent on the presence of one particular element or on the presence or absence of a solvent.
  3. It explains the long-accepted basic properties of metallic oxides and acidic properties of non-metallic oxides.
  4. This theory explains many reactions such as gas-phase high temperature and non-solvent reaction as the neutralisation process.
  • Problem 1:
    Explain - NH3 behaves as a base but BF3 as an Acid.
  • Answer:
    In NH3, the central N atom has 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 contains six electrons in the central atom thus it behaves as an acid.
F3B + : NH3 F3BNH3
Acid Base Adduct
  • Problem 2:
    Explain why tri-covalent phosphorus compounds can serve both as Lewis acids and also as bases?
  • Answer:
    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
Acid Base Adduct
    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.
  • Problem 3:
    Bi-positive tin can function both as a Lewis acid and a Lewis base. Explain?
  • Answer:
    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 the Lewis base behaviour of SnCl2.
  • Problem 4:
    Can SiCl4 and SnCl4 function as Lewis acids?
  • Answer:
    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.
  • Problem 5:
    Arrange these oxides in order of their acidic nature: N2O5, As2O3, Na2O, MgO.
  • Answer:
    Acidic oxides react with water to give oxoacids. The higher the oxidation number and the higher the electronegativity the greater 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: Na2OMgOAs2O3N2O5 Oxidation number of Oxide forming element: +1+2+3+5.
  • Problem 6:
    Bisulphate ion can be viewed both as an acid and a base. Explain.
  • Answer:
    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.
  • Problem 7:
    SO3 Behaves as acid and H2O as a base. Explain.
  • Answer:
    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.

Lewis Concept of Acids and Bases with Classification, Bronsted, Arrhenius and solvent System Concept Vs Lewis Concept and related problems with solutions

Inorganic Chemistry

[Inorganic chemistry][column1]

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