Acid and Base

Concept of acid and base

1. Arrhenius concept

Depending upon the types of ions produced by a substance in water, Arrhenius proposed a modern concept of acids and bases. According to it, an acid is a compound that releases hydrogen ions in water.
HCl, HNO3, etc. are acids because they release hydrogen ions in water.
According to this concept, common properties of acids i.e. having a sour taste, the ability to turn blue litmus into red, the ability to neutralize with base, etc are due to the presence of H⁺ ions in their aqueous solution.

Similarly, a base is a compound that releases hydroxide ion in water. NaOH, KOH, etc. are bases because they release OH^– ion in water.
According to this concept, the common properties of bases i.e. having a bitter taste, the ability to turn red litmus into blue, the ability to neutralize with acids, etc. are due to the presence of OH^– ion in the aqueous solution.

Limitations

1. No free existence of ions: They immediately associate with water to form other types of ions.
2. Limited scope: This concept is applicable when the solvent taken is water.
3. There are many acids and bases without H+ and OH- ions in their original formula like CO₂, MgO, etc.

2. Bronsted-Lowry concept

According to this concept, an acid is any substance (molecule, cation, or anion) that can donate a proton.

HCl + H_{2}O\ \rightarrow\ H_{3}O^{+} + Cl^{-}\\ HSO_{4}^{-} + H_{2}O\ \rightarrow\ H_{3}O^{+} + SO_{4}^{--}

HCl and HSO₄^– are Bronsted acids because they donate a proton.

Similarly, a base is any substance (molecule, cation, or anion) that can accept a proton.

NH_{3} + HCl\ \rightarrow\ NH_{4}^{+} + Cl^{-}\\ CN^{-} + H^{+}\ \rightarrow\ HCN

NH₃ and CN^– are Bronsted bases because they accept a proton.

This concept is superior to Arrhenius’s concept due to the following reasons:

i. Much wider scope: Not only the neutral molecule but also the ions are defined as acids and bases according to this concept. So it has a wide concept.

ii. Aqueous solution is not always an essential condition to define acidic or basic character.
NH₃(g) + HCl(g) → NH₄Cl
Here, NH₃ is an acceptor of the proton and acts as the Bronsted base and HCl is a donor of the proton and acts as Bronsted acid but the reaction can be carried out in the gases phase.

iii. Release of OH^– ion is not always necessary to qualify a base.
NH₃ + H⁺ → NH₄⁺
Here, NH₃ acts as a Bronsted base by accepting a proton.

Limitations

i. The concept is unable to explain the reaction of non-protonic acids like CO₂, SO₃, etc. with bases like CaO, MgO, etc.
ii. The concept could not explain the acidic behaviour of non-protonic substances like BF₃, AlCl₃, FeCl₃, etc.

Conjugate acid-base pair

A pair of acids and bases which differ from each other by a proton is called a conjugate acid-base pair.

Auto Ionization of water

H₂O + H₂O ⇌ H₃O⁺ + OH^–
The self-interaction of the water molecules to produce hydronium and hydroxide ions is called the autoionization of water. Here, one molecule of water is a donor of proton and acts as Bronsted acid whereas another molecule of water is an acceptor of proton and acts as Bronsted base.

Water as amphoteric substance

The amphoteric nature of water can be explained by taking two Bronsted acid-base reactions:

H₂O + NH₃ ⇌ NH4⁺ + OH^–
Here, H₂O is the donor of proton and acts as Bronsted acid.

H₂O + HCl ⇌ H₃O⁺ + Cl^–
Here, H₂O is an acceptor of the proton and acts as a Bronsted base.

Relative strengths of Bronsted acid and base

The strength of bronsted acid and bronsted base depends upon the tendency to donate and to accept proton respectively. Greater the tendency for it, stronger is the acid or base. Let us consider a bronsted acid-base reaction:

Since the process shifts almost towards the forward direction, HCl must have a strong tendency to donate a proton and it acts as strong acid. At the same time, H₂O must have a strong tendency to accept a proton and act as a strong base. On the other hand, since the process has very little tendency to shift towards the backward direction, both the species on the right-hand side are weak.

Conclusions
i. Strong acid has weak conjugate base and vice versa.
ii. Strong base has weak conjugate acid and vice versa.

3. Lewis concept

This is the electronic concept. According to this concept, an acid is any molecular substance that can accept a pair of electrons.
According to this concept, the following species can act as Lewis acid:
i. Simple cations (Na⁺, Ag⁺, H⁺)
ii. Neutral molecule containing the central atom of incomplete octet (AlCl₃, BF₃, etc.)
iii. Molecules containing atoms of different electronegativity linked to each other by multiple bonds (CO₂, SO₃, etc.).

Similarly, a base is any molecular substance that can donate a pair of electrons. According to this concept the following species can act as Lewis base:
i. Simple anions (SO₄^–, NO₃^–, Cl^–, etc.)
ii. Neutral molecules containing at least one lone pair of electrons (NH₃, H₂O, R-OH, R-NH₂, etc.)

Limitations

1. It is unable to explain the strength of acids and bases.
2. According to this theory, there is the formation of a coordinate bond while acid reacts with the base but no such bond can be observed for some well-known acid-base reactions.
   HCl + NaOH → NaCl + H2O
   Here, an ionic bond is formed.
3. This concept could not explain the catalytic activity of acid.