Lewis concept of acids and bases

Lewis concept of acids and bases
Lewis concept
In 1923, the American chemist Gilbert Newton Lewis proposed generalized definitions for acids that do not restrict them to compounds containing hydrogen but in terms of the structure of acids and bases. Lewis defined an acid as any molecular or ionic species that contains an atom capable of sharing a pair of electrons furnished by a base. In the simplest terms, a Lewis acid is an electron acceptor.
    Lewis base was to be considered according to Lewis as an electron pair donor and Lewis acid as an electron pair acceptor.
    Lewis 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 the lewis, the neutralization reaction is simply the formation of a coordinate bond between an acid and a base.
    The neutralization product, termed as coordinate complex or adduct, may be either non-ionizable or may undergo dissociation or condensation reaction depending on its stability.
H⁺ + : NH₃ H ← : NH₃]⁺
Acid Base Adduct
F₃B + :NH₃ ⇄ F₃B ← :NH₃

Lewis vs Bronsted-Lawry acids and bases

    Bronsted-Lawery defined, a base is any molecules or ion, which accepts a proton. Lewis defined an acid as any molecular or ionic species that contains an atom capable of sharing a pair of electrons furnished by a base. So that during neutralization reaction a coordinate covalent bond is formed between the donor atom of the base and the proton.
H⁺ + :NH₃ ⇄ [H ← :NH₃]⁺
    Thus all Bronsted-Lawery bases are bases in Lewis sense also.

Arrhenius vs Lewis concept of acids and bases

    Arrhenius defined, 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 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.

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

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.

The simple cation of an atom

    Theoretically, all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis base are,
Lewis acid + Lewis BaseAdduct

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⁺²ㄑFe⁺³ ( Positive charge increases from +2 to +3) K⁺ㄑNa⁺ (On moving from bottom to top in a group) Li⁺ㄑBe⁺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 sulfur or carbon atom. As a result, the electron density of π-electrons is displaced away from carbon to sulfur atoms which are less electronegative than oxygen, towards the O-atom.
    Thus carbon and sulfur are electron deficient able to accept the electron pair from a Lewis base such as OH⁻ ions to form a dative bond.

Element with an electron sextet

    Oxygen and sulfur 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 sulfur is the acid-base reaction.
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 is involved.
  2. Lewis's 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.
  3. It explains the long-accepted basic properties of metallic oxides and the acidic properties of non-metallic oxides.
  4. This theory explains many reactions such as gas-phase high temperature and non-solvent reaction as the neutralization process.

Questions answers

Question
    Explain - ammonia behaves as a base but boron trifluride as an Acid.
Answer
    In ammonia, the central nitrogen atom has lone pair of electrons This lone pair coordinate to the empty orbital of an atom, which is termed as a base according to Lewis.
    The compounds with less than an octet for the central atom are Lewis acids. In BF₃, B in BF₃ contains six electrons in the central atom thus it behaves as an acid.
F₃B + : NH₃ F₃B←NH₃
Acid Base Adduct
Question
    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 with 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-suborbital which can receive back donated electrons from electron reach low oxidation state of a metal ion. In this latter case, the tri-covalent phosphorus compound serves as a Lewis acid.
Question
    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)PtCl(SnCl₃)], [RuCl(SnCl₃)₂]⁻²
    These are examples of the Lewis base behavior of SnCl₂.
Question
    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 hybridization 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.
Question
    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
Na₂OㄑMgOㄑAs₂O₃ㄑN₂O₅

Oxidation number of oxide-forming element
+1, +2, +3, +5.
Question
    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.
Question
    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 and classification of acids and bases, Bronsted, Arrhenius and solvent system concept Vs Lewis concept with questions answers

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