Definition of Lewis Acid and Base
Lewis acid base concept or theory of acid and base in 1923 explains the acids bases properties in terms of electronic structure with the formation of the coordinate covalent bond in chemistry. According to American chemist Gilbert Newton Lewis definition, an acid is any species (molecule, radical, or ion) that can accept an electron pair to form a coordinate covalent bond and a base is any species that can donate an electron pair to the formation of a coordinate covalent bond. In learning chemistry, the Lewis acids and bases concept does not restrict them to concentration hydrogen ion or pH scale but in terms of the structure of atoms, ions, or molecular species.
In the Lewis acid base neutralization reaction, an acid act as an electron acceptor and base considered as an electron pair donor. The hard soft acid base theory explaining the stability of the formation complex.
Electronic Theory of Lewis Acids and Bases
Lewis acid, therefore, equilibrium reaction with a base leading to a coordinate link from the base to the acid. Thus an acid must have a vacant orbital into which an electron pair donated by a base can be accommodated. A Lewis base is a substance which has at least one unshared (lone) pair of electrons. According to the lewis, the neutralization reaction is simply for the formation of the co-ordinate bond between an acid and a base. The neutralization product, termed as acid base complex or adduct, may be either non-ionizable or may undergo ionization depending on its stability.
Bronsted and Lewis Acids and Bases
According to the protonic definition by Bronsted-Lawery, a base is any molecules or ion, which accepts a proton. Lewis defined an acid as any molecular or polar species that contains an atom capable of sharing a pair of electrons furnished by a base. Therefore, during the neutralization reaction, a coordinate covalent bond is formed between the donor atom of the base and the proton or hydrogen ion (acids). Hence all Bronsted-Lawery bases are bases in Lewis sense also.
Question: What is the conjugate acid of bisulfate?
Answer: Bisulphate ion is HSO3–. It may lose hydrogen ion to give the conjugate base SO3-2, thus behaving as acids. Again it may use hydrogen ion to give the conjugate acid, thus showing its base character.
Arrhenius and Lewis Acids and Bases
Arrhenius defined, acids when dissolved in water dissociate into hydrogen ions and anions. Therefore, hydrogen ion (H+) released from nitric acid, and sulfuric acid (responsible for acid rain) acts as acids. Since the hydrogen ion has the affinity to receive an electron pair from the base. Thus all the Arrhenius acids are also acids under the Lewis definition.
Acid and Base in Solution
- The solvent system concept recognizes that ammonium chloride should behave as an acid in liquid ammonia, alcohol, or acetic acid solution. But an ammoniated hydrogen ion (acids) can accept a lone pair of an electron from the base. Therefore an ammonium ion in liquid ammonia solution 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.
Question: What type of chemical bond formed, when sulfur trioxide reacts with water?
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 Acids and Bases in Chemistry
A wide variety of chemical compounds are recognized as acids and bases according to Lewis in chemistry. Any Lewis acid must contain at least one empty orbital in the valence shell electron configuration.
Incomplete Octet of the Central Atom
Examples of these acids are electron-deficient molecules such as alkyl and halides of beryllium, boron, and aluminum. Some reactions of these types of Lewis acid with Lewis base are,
F3B + O(C2H2)2 ⇄ F3B ← :O(C2H5)2
Cl3Al + NC5H5 ⇄ Cl3Al ← :NC5H5
Me3B + N2H4 ⇄ Me3B ← :N2H4
Question: What happens when BF3 reacts with NH3?
Answer: In ammonia, the central nitrogen atom has a 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, boron contains six electrons in the central atom thus it behaves as an acid.
Elements having vacant d orbital
The central atom of the halides such as SiX4, GeX4, TiCl4, SnX4, PX3, PF5, SF4, SeF4, TeCl4, etc. have vacant d-subshell. These chemical substances which can accept an electron charge pair from the Lewis base to accommodate in their vacant d-subshell and form adducts with a number of halide ions and oxygen, nitrogen, and sulfur-containing organic hydrocarbon. 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 Lewis acid base adduct (intermediate) with H2O.
Problem: Why PCl3 can act both as lewis acid and lewis base?
Solution: PCl3 has a lone pair of electrons in phosphorus. This lone pair may coordinate with a metal ion thus the compound has properties of Lewis base.
Ni + 4PCl3 ⇄ [Ni(PCl3)4]0
Again the quantum shell of phosphorus has provision for d sub-orbital which can receive back donated electrons from electron reach low oxidizing agents metal ion. In this latter case, the tri-covalent phosphorus compound serves as a Lewis acid.
Cation is Formed when an Atom
Theoretically, all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis base are,
Ag+ + 2(:NH3) ⇄ [H3N: → Ag ← :NH3]+
Cu+2 + 2(:NH3) ⇄ [Cu(NH3)4]+2
Co+3 + 6(:OH2) ⇄ [Co(OH2)6]+3
Li+ +: OHCH3 ⇄ [Li ← :OHCH3]
Strength of Acid and Base
The Lewis acid strength or coordinating ability of the simple cations are increased with-
- An increase in the positive charge carried by the cation.
- An increase in the charge of the nucleus or mass number of an atom in any period of the periodic table.
- A decrease in ionic radius.
- A decrease in the number of shielding electron shells.
- Hydrogen bonding also affects the acid base strength.
Evidently, the Lewis 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. Fe+2ㄑFe+3 ( oxidizing properties decreases), 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 polarization between atoms of dissimilar electronegativity. For examples of this class of Lewis acid base like CO2, SO2, SO3, borax, hydrogen peroxide, boric acid, etc in these compounds, the oxygen atom is more electronegative than the central atom. As a result, the density of π-electrons is displaced away from carbon to sulfur atoms which are less electronegative than oxygen, towards the O-atom. Therefore, carbon and sulfur are electron deficient chemical compounds of our environment able to accept the electron pair from Lewis bases such as OH– ions to form a dative bond with the very low bond energy.
Lewis Acidic Nature of Oxides in Chemistry
Acidic oxides react with water to give oxoacids in redox reactions. The higher the oxidation number and the higher the electronegativity the greater the central chemical 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: Na2OㄑMgOㄑAs2O3ㄑN2O5 and oxidation number of oxide-forming element = +1, +2, +3, +5
Chemical Properties of Acid and Base
- Lewis acids bases concept includes those chemical reactions in which no hydrogen atom or ion is involved.
- Lewis’s concept is more general than the Bronsted – Lowry concept or protonic concept in that acid-base properties are not dependent on the presence of one particular element or on the presence or absence of a solvent.
- It explains the long-accepted base properties of metallic crystalline oxides and the acidic properties of non-metallic oxides.
- Lewis concept or theory explains many acid base reactions in the gas phase in high temperatures and non-solvent reactions as the neutralization process.