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1.3.1 LEWIS ACIDS AND BASES

According to G N Lewis (1923):

1. A Lewis base is a neutral molecule or an ion that can donate a pair of electrons to other molecules or ions. For example, ammonia (NH₃), water (H₂O), hydroxide (OH⁻) or chloride (Cl⁻) ions.
2. An Lewis acid is a molecule or an ion that can accept a pair of electrons from an electron pair donor (Lewis base). Examples include hydronium ion (H⁺), ammonium ion (NH₄⁺), boron trifluoride (BF₃), etc.
3. Neutralization is the process of donation of a pair of electrons from a Lewis base to an Lewis acid. In the process, a coordinate bond is formed. Hence, in addition to the usual Brönsted acid–base neutralization reactions, the Lewis definition considers the formation of a coordinate bond also as a neutralization reaction:

    BF₃    +     NH₃     →   H­₃N⁺–BF₃ ^–
Lewis acid + Lewis base  → Coordinated adduct

  1.3.1.1 Lewis Acids

Lewis acids are defined as those compounds or ions, which cannot be classified as acids on the basis of Arrhenius or Brönsted definition, but can accept a pair of electrons from other molecules or ions. Thus, the following species are Lewis acids:

1. All cations, as they can combine with electron pairs forming a coordinate bond:

Ag⁺ + 2CN^– → [Ag(CN)­₂] ^–
Ni²⁺ + 4NH₃ → [Ni(NH₃)₄]²⁺
Fe³⁺ + 6CNS^– → [Fe(CNS)­₆]^{3 –}
Al³⁺ + 4OH^– → [Al(OH)­₄] ^–
Co³⁺ + 6NH₃ → [Co(NH₃)­₆]³⁺

2. Electron deficient compounds, in which the central atom does not have a complete octet of electrons. For example, boron trifluoride (BF₃) or aluminium chloride (AlCl₃). In these molecules, the boron atom or aluminium atom has only six electrons and is two electrons short of a complete octet in its outermost valence shell. These compounds are termed electron deficient compounds. . In order to complete their octet and attain stability, these atoms can accept a pair of electrons from other electron pair donors (bases) like ammonia or chloride ion, complete their octet and form a stable species, called adducts which are coordination compounds:

BF₃    +     NH₃     →   [H­₃N⁺–BF₃ ^–]
AlCl₃  +     Cl ^–        →    [AlCl₄] ^–

3. Covalent compounds in which the central atom has a complete octet of electrons, but can expand its valence shell by accepting one or more pairs of electrons. This is called the expansion of valence shell to more than an octet. Covalent halides of element of third to sixth period in group 14 to 17 like SiCl₄, SnCl₂, SnCl₄, PCl₅, AsF₅, SbF₅, SF₄, BrF₃, etc. are Lewis acids, as they can accept one or more electron pairs from say, the halide ions, and form coordinated complexes:

SnCl₄ + 2Cl ^–  → [SnCl₆]^{2 –}
PCl₅ + Cl ^–       → [PCl₆] ^–
AsF₅  +  F ^–     → [AsF₆] ^–
SbF₅ + Cl ^–     → [SbClF₅] ^–
IF₅  +  F ^–         → [IF₆] ^–

4. Multiple bonded covalent compounds SO₂, SO₃, CO or CO₂, which can form a covalent bond with anions like halide, hydroxide or other Lewis bases:

4. SO₂ + OH ^–  → [SO₂(OH)] ^–
   SO₃ + F ^–  → [SO₃F] ^–
   CO₂ + OH ^–  → [CO₂(OH)] ^–
   CO + Cl₂  → COCl₂

1.3.1.2 Lewis Bases

Lewis bases are defined as those compounds or ions, which cannot be classified as a base on the basis of Arrhenius or Brönsted definition, but contain lone pair(s) of electrons that can be donated to Lewis acids, and form a coordinate or covalent bond.

Due to the presence of lone pairs of electrons, almost all the Lewis bases are Brönsted bases.

Lewis bases can be classified as follows:

1. Molecules having an unshared or lone pair of electrons, i.e., water, ammonia, ethers (R₂O), alcohols (ROH), amines (RNH₂, R₂NH, R₃N), etc. Lewis base strength increases with the increase in the availability of the electron pair.
2. All anions i.e., OH ^–, S^{2 –}, Cl ^– or NO₂ ^– as all anions have (a) an atom with free electron pairs and (b) a negative charge density. Lewis base strength increases with increase in charge density on the atom having the lone pair making it easier for donation.
3. Compounds having multiple-bonded atoms like carbon monoxide, nitric oxide, alkenes, alkynes; (multiple bonded molecules) and conjugated organic compounds having multiple bonds, e.g., benzene or allyl radicals. These species can combine with transition metals and their ions and form complexes containing coordinate bonds e.g., [Ni(CO)₄], [Fe(H₂O)₅(NO)]²⁺, [Ag(CH₂=CH₂)]⁺, [Cr(C₆H₆)₂], etc.

1.3.2 LUX–FLOOD  DEFINITION                                                       

The definition proposed by Lux (1939) and extended by Flood (1947) considers the acid–base reaction on the basis transfer of oxide ions. Acid is a substance that can accept oxide ions (O^2–) whereas bases are the oxide ion donors.

Thus the reactions

 CaO + SiO₂ → CaSiO₃
 ZnO + SO₃ → ZnSO₄

involve basic oxides (CaO, ZnO) and acidic oxides (SiO₂, SO₃) to form salts (CaSiO₃, ZnSO₄).

The acids need not always be oxides as can be seen in the interaction of pyrosulphates with oxides of titanium, niobium or tantalum at 1100 K:

TiO₂  +         Na­₂S₂O₇  → TiOSO₄ + Na₂SO₄ 
 Base           Acid                    Salts

Substances are called amphoteric if they tend to both lose and gain oxide ions:

  ZnO ⇌ Zn²⁺  + O^2– (Base) ; ZnO + O^2– ⇌ ZnO₂^2- (Acid)
 Al₂O₃ ⇌ Al³⁺ + 3O^2– (Base) ; Al₂O₃ + O^2– ⇌ 2AlO₂^–  (Acid)

The Lux−Flood definition  can be extended to include any anion transfer like halide or sulphide. It is applicable to systems where no solvents involved like molten melts or non-protonic solvents.

1.3.3 USANOVICH DEFINITION

Usanovich removed the condition that bases should donate a pair of electrons as such and extended the Lewis definition by proposing that  any number of electrons can be transferred in an acid−base  reaction. According to Usanovich’s definition, acid is a chemical species that can accept one or more electrons (not necessarily electron pairs), whereas base is the chemical species that can donate or release any number of electrons.

 Therefore, in addition to Arrhenius, Brönsted and Lewis acids, the Usanovich’s acids include all the oxidizing agents (electron acceptors), whereas bases include all the reducing agent (electron donors). When a species accepts electrons, it is said to undergo reduction, and acts as an oxidizing agent. Thus, in reactions

Ag⁺ + e^– ⇌ Ag 
S + 2e^– ⇌ S ^{2 –} 
Fe³⁺  + 3e ^–  ⇌ Fe²⁺

In these examples, silver ions, sulphur and iron(III) ions are accepting electrons and are getting reduced. However, according to Usanovich definition, these reactions show the acidic behavior of these ions.

Similarly, when a species loses electrons, it undergoes oxidation and acts as a reducing agent. Thus, hydrogen and metals are reducing agents – they  lose electrons and undergo oxidation:

 H₂ ⇌ 2H⁺ + 2e^–
 Na ⇌ Na⁺ + e ^–
Mg ⇌ Mg²⁺ + 2e ^–

 But Usanovich considers such species as  bases as they lose electrons in the reactions and are electron donors.

However, in some cases, there is not much difference between the Usanovich definition and the Lewis definition. For example, the reaction between pyridine and oxygen to form pyridine-N-oxide is an oxidation reaction of pyridine as well as the adduct formation in which pyridine nitrogen atom donates an electron pair to oxygen atom:

2C₅H₅N + O₂ → [C₅H₅N→O] or [C₅H₅N⁺–O^–[C₅H₅N→O]

1.3.4 CADEY – ELSEY DEFINITION

Cadey – Elsey definition of acids and bases is based upon the auto-ionization of solvents. In liquid state, all the ionizing solvents undergo auto-ionization, i.e., dissociate into their ions:

 Solvent (l) ⇌ Cation (Acid)  + Anion (Base)
2H₂O         ⇌       H₃O⁺      +             OH ^–
2NH­₃(l)      ⇌       NH₄⁺      +             NH₂ ^–
2SO₂(l)      ⇌        SO²⁺       +             SO₃ ^{2 –}
CH₃CO₂H(l) ⇌    CH₃CO₂H₂⁺  +     CH­₃CO₂ ^–

In an ionizing solvent like water, liquid ammonia, liquid sulphur dioxide or acetic acid, acids are defined as those species that increase the concentration of the cation of the solvent , whereas bases are defined as those species that increase the concentration of the anion of the solvent. Thus, in water as solvent or in aqueous solutions, all the species that increase the concentration of hydronium ions (H₃O⁺) are acids, whereas the species that increase the concentration of hydroxide ions (OH ^–) are bases.

However, in liquid ammonia as the solvent, the acids are compounds like ammonium chloride (NH₄Cl), ammonium nitrate (NH₄NO₃) and ammonium acetate (CH₃COONH₄) which ionize to form ammonium ions and increase the concentration of the cation of the solvent:

 NH₄Cl  ⇌ NH₄⁺  + Cl^–
NH₄NO₃ ⇌ NH₄⁺  + NO₃^–
CH₃COONH₄  ⇌  NH₄⁺ + CH₃COO ^–   

 Similarly, in liquid ammonia as solvent, bases are compounds like sodamide (NaNH₂) or potassium amide (KNH₂), which ionize to form NH₂^–, the amide anion:

NaNH₂ ⇌ Na⁺ + NH₂^–

The advantage of such a definition is that all the ionizing systems AB can be treated analogously. The ionic product of the solvent can be defined as K_AB

 K_AB = [A⁺][B ^–]

And, a pA  scale, analogous to the pH scale, can be set for each solvent where the neutral point when [A⁺] = [B ^–] will be given by –(1/2)K_AB,

Limitations: The limitations of the solvent system definition are the following.

1. Emphasis is only on the auto-ionization os the solvent; other physical characteristics are comletely ignored.
2. Due to low dielectric constants of many of these solvents, existence of ions in the solvent are is not conceivable.
3. For many solvents like POCl₃, auto-ionization is stretched too far.
4. For many reactions, ionic reactions are too emphasized where other mechanisms are sufficient to explain the reaction.

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