Lewis acid and base definition
Lewis acid base concept or theory of acids and bases in 1923 explains the acid base properties in terms of electronic structure with the formation of the coordinate covalent chemical bond in chemistry. According to American chemist Gilbert Newton Lewis, an acid is any species (molecule, radical, or ion) that can accept an electron pair to form a coordinate covalent bond. A base is any species that can donate an electron pair to the coordinate covalent bond. In learning chemistry, the Lewis acids and bases concept does not restrict to the concentration of hydrogen ions or pH scale of the aqueous solution. It is based on the electronic configuration or structure of atoms, ions, or molecular species.
Lewis acid-base theory
In acid base chemical equilibrium reaction, a base leading to form coordinate link to the acid. An acid must have a vacant orbital in which the electron pair donated by a base can be accommodated. A Lewis base is a substance that has at least one unshared (lone) pair of electrons. According to Lewis theory, the neutralization reaction is simply for the formation of the co-ordinate bond between an acid and a base. The neutralization product is called acid base complex or adduct. It is either non-ionizable or may undergo ionization depending on its stability.
Lewis acids and bases examples
In the Lewis acid base neutralization reaction, an acid act as an electron acceptor, and the base is considered as an electron pair donor. In the above picture, we provide some examples of Lewis acids and bases with acid base complexes. The stability of such complexes explains by hard soft acid base theory.
Arrhenius acid base theory
Arrhenius defined, acids when dissolved in water dissociate into hydrogen ions and anions. Therefore, hydrogen ion (H+) are released from nitric acid, and sulfuric acid acts as an acid. Since the hydrogen ion has the electron affinity to receive an electron pair from the base. Thus all the Arrhenius acids are also acids under the Lewis definition.
Bronsted Lowry acid base
According to the protonic definition by Bronsted-Lowry, a base is any molecule or ion, which accepts a proton. Lewis defined an acid as any molecular or polar molecule that contains an atom capable of sharing a pair of electrons to a base. 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-Lowry bases are bases in the Lewis sense also.
Question: What is the conjugate acid base pair 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.
Solvent system concept
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 an acid according to Lewis sense. The solvent system considers amide ion as a base in liquid ammonia. The amide ion is capable of donating a lone pair to an acid, hence this is a base according to Lewis theory.
Question: What type of chemical bonding formed, when sulfur trioxide reacts with water?
Answer: Sulfur trioxide (SO3) like boron trifluoride (BF3) has less than an octet and will be termed as an acid according to Lewis concept. But in H2O oxygen atoms contain lone pairs to donate the Lewis acid.
Classification of Lewis acid and base
A wide variety of chemical compounds are recognized as acid and base according to Lewis theory in chemistry. According to Lewis theory, acids and bases are classified as,
- incomplete octet of the central atom
- elements having vacant d orbital
- cations of an atom
Incomplete octet of the central atom
|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 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 with 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?
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 of an atom
|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]|
Lewis acids and bases strength
The Lewis acid strength or coordinating ability of the simple cations is increased with an increase in the positive charge carried by the cation. With an increase in the charge on the nucleus or the atomic number of the metals in any period of the periodic table, acid strength is increasing.
It decreases by decreasing ionic radius the number of shielding electron shells. Hydrogen bonding also affects acid-base strength. For example, 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).
Example of Lewis acid
Molecules having multiple electric polarization between atoms of dissimilar electronegativity like carbon dioxide, sulfur dioxide, sulfur trioxide, borax, hydrogen peroxide, boric acid are examples of lewis acids. In these compounds, the oxygen atom is more electronegative than the central atom. As a result, π-electron density is displaced away from the central atom. Therefore, these are electron deficient chemical compounds able to accept the electron pair from Lewis bases. They form a dative bond with very low bond energy.
Acidic nature of oxides
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 react with water to give the oxoacid. Of all these oxoacids nitrogen has the highest oxidation number and electronegativity. The order of acidic nature: Na2OㄑMgOㄑAs2O3ㄑN2O5 and oxidation number of oxide-forming element = +1, +2, +3, +5
Utility of Lewis acid base concept
- Lewis acids bases concept includes those chemical reactions in which no hydrogen atom or ion is involved.
- The 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 solid oxides and the acidic properties of non-metallic oxides.
- The Lewis concept or theory explains many acid base reactions in the gas phase in high temperatures and non-solvent reactions as the neutralization process.