Strength Of Acids And Bases

Strength Of Acids And Bases

There are various theories regarding acids & bases

(i) Arrhenius Theory  

According to Arrhenius theory, substances producing H+ ions in solution are acids and those producing OH- ions in solution are bases. Therefore, substances like H2O, HCl, H2SO4, CH3COOH etc. are acids and the ones like NH₄OH, NaOH, KOH, H₂O etc. are bases.

(ii) Bronsted−Lowry Theory

In 1923, Bronsted and Lowry independently defined acids as proton donors, and bases as proton acceptors. For aqueous solutions the definition does not vary much for acids from the Arrhenius theory but it widens the scope of bases. In this, the bases need not contain OH- ions and simply have to accept protons. So ions like Cl-, CH₃COO-, Br- etc. which do not contain OH- ions can be considered as bases under this definition.

Levelling Solvents: Whenever an acid is dissolved in water, it acts as an acid only if the solvent acts as a base. That is, if we dissolve acids like HCl, HNO₃ etc in water, their acidic strength is almost the same since water acts as a base for both these acids. Infact, it is known that all strong acids show equal acidic strength when dissolved in water. This is because, water acts as a base to all these acids and thus forces them to donate almost the same amount of protons irrespective of their chemical nature. Since water levels the acidic strength of strong acids, it is referred to as a levelling solvent. In order to measure the strength of strong acids, they are dissolved in glacial acetic acid and the amount of protons is measured by conductometry. It is found that the strength of acids varies as:

HClO₄ > HBr > H₂SO4 > HCl > HNO₃

(iii) Lewis Theory

Lewis developed a definition of acids and bases that did not depend on the presence of protons nor involve reactions with the solvent. He defined acids as materials which accept electron pairs, and bases as substances which donate electron pairs. Thus a proton is Lewis acid and ammonia is Lewis base since, the lone pair of electrons on the nitrogen atom can be donated to a proton:

H+ + :NH₃ →  [H ← :NH₃]+

Ag+ + 2 NH₃  → [H₃N: → Ag ← :NH₃]+

Lewis acid Lewis base Acid-base adduct

Conditions to be a Lewis Acid:

(i) Compounds whose central atoms have an incomplete octet e.g. BF₃, AlCl_3, GaCl₃ etc.

(ii) Compounds in which the central atom have empty d-orbitals and may acquire more than an octet of valence electrons.e.g.

SiF₄ + 2F− →  SiF₆²⁻

Other examples are : PF₃, SF₄, SeF₄, TeCl₄.

(iii) All simple cations : Na+, Ag⁺, Cu^2+, Al³⁺, Fe³⁺ , Mg²⁺, Ca²⁺ etc.

Conditions to be a Lewis Base:

(i) All simple negative ions e.g. Cl− , F−, O₂−, SO^-2₄ etc.

(ii) Molecules with unshared pair of electrons: H₂O, NH3 etc.

(iii) Multiple bonded compounds which form co-ordination compounds with transition-metals,e.g., CO, NO, Ethylene, Acetylene etc.

Amphiprotic Species

Water can either gain or lose a proton and thus can behave as an acid as well as a base. Such species are called amphiprotic species. Similarly many other molecules and ions can either gain or lose a proton.

Examples:

Acid Base Acid Base

HS− + OH− H2O + S2−

HBr + HS− H2S + Br−

HSO₄₋ + OH− H2O + SO₄²⁻

HClO₄ + HSO₄− H₂SO₄ + ClO⁴⁻

The oxides and hydroxides of metals near the boundary between metals and non-metals are amphiprotic and can react either as acids or as bases.

Hard and Soft Acids and Bases (HSAB Principle)

Hardness is measured as the property of retaining valence electrons very strongly. Lewis acids and bases can be classified as hard and soft acids and bases as follows:

Hard Acid is that in which the atom, which is accepting electrons, is smaller in size and has a high positive charge.

Soft Acid is that in which the atom, which is accepting electrons, is bigger in size, has low positive charge and the electrons in it can be easily polarised.

Hard Base is that in which the electron donating atom is small and has high electronegativity e.g. F-, NH₃, H₂O, OH- etc.

Soft Base is that in which the electron donating atom is bigger and has low electronegativity e.g. I-, PH₃, (CH₃)₃P etc.

A hard acid prefer to bind to hard bases to form ionic bond and the soft acids prefer to bind to soft bases to form mainly covalent bonds.