Acid and Base

Theories of Acids and Bases:

-There have been many attempts to define the difference between acids and bases. The three main theories in use today are:

Arrhenius's concept of acids and bases
Bronsted–Lowry's theory of acids and bases
Lewis's theory

ARRHENIUS'S Concept of Acids and Bases:

-Arrhenius defined acids and bases as follows:

An acid is a substance that produces hydrogen ions when dissolved in water.
A base is a substance that produces hydroxyl ions when dissolved in water.

Examples of acids: $H Cl, H_{2} S O_{4}, H N O_{3}$

$H Cl + H_{2} O ⇋ H_{3} O^{+} + C l^{-}$

Examples of bases: $N a O H, K O H, e t c$ .

$N a O H + H_{2} O ⇋ N a^{+} + O H^{-}$

Limitations of Arrhenius Theory of Acids and Bases:

It recognises the dissociation of acids and bases in an aqueous medium only.
It restricts acids to merely hydrogen-containing compounds and bases to merely hydroxide-containing compounds.

According to Arrhenius concept, carbon dioxide, sulphur dioxide, etc., are not regarded as acids and $N H_{3}, C a O, M g O, e t c .$ are not regarded as bases.

Bronsted–Lowry’s Concept of Acids and Bases:

-Bronsted and Lowry proposed a concept of acids and bases, which are independent of solvent.

-According to this concept, acids and bases are defined as follows:

Bronsted Acid:

-An acid is a species (a molecule, a cation (or) an anion) capable of donating one or more protons to any other substance called Bronsted-Lowry acid.

-Bronsted acids dissociate to release protons, resulting in a higher concentration of H⁺ ions in the solution.

$A c i d ⇌ P ro t o n + C o nj ug a t e ba se$

-Strong Bronsted-Lowry acids are those that have a strong inclination to donate a proton but have a weak conjugate base.

-Weak Bronsted-Lowry acids have a slight tendency to give a proton, while their conjugate base is strong.

Examples for acids: $H Cl, H_{2} O, H_{2} S O_{4}, H N O_{3}$

$H_{2} O ⟶ H^{+} + O H^{-}$

Bronsted Base:

-A base is a species (a molecule, a cation (or) an anion) capable of accepting one or more protons from an acid; such a base is called the Bronsted-Lowry base.

$B a se + P ro t o n ⇌ C o nj ug a t e a c i d$

Examples for bases: $N H_{3}, O H^{-}, C N^{-}$ $N H_{3} + H^{+} ⟶ N H_{4}^{+}$

Therefore, an acid is a proton donor, and a base is a proton acceptor.

Limitations of Bronsted-Lowry's Concept:

Cannot explain the acidic nature of $C O_{2}, S O_{2}, S O_{3}$
Cannot explain the basic nature of $C a O, M g O, B a O$
Cannot explain how chemicals that lack hydrogen, such as $B F_{3} an d A lC l_{3}$ , display acidic characteristics.

Lewis Concept of Acids and Bases:

Lewis acid:

-A species with an empty orbital and hence the ability to take an electron pair.

-Lewis acids have electrophilic properties.

- $C u^{2 +}, B F_{3}, an d F e^{3 +}$ are examples of Lewis acids.

-A Lewis acid absorbs an electron pair from a Lewis base, resulting in the formation of a coordinate covalent bond.

Lewis base:

-A Lewis base is a species that has a single pair of electrons and hence can operate as an electron donor.

-Lewis bases have nucleophilic properties.

- $F^{-}, N H_{3}, an d C_{2} H_{4} (e t h y l e n e)$ are examples of Lewis bases.

The hydrogen atom is not included in this theory's description of acids and bases.
This notion has the advantage of allowing various substances to be classified as acids or bases. However, it provides little information about the acid and base's strength.
One of the theory's flaws is that it doesn't account for acid-base reactions that don't involve the development of a coordinate covalent bond.

pH of Acids and Bases:

To find the numeric value of the acidity or basicity level of a substance, the pH scale (pH stands for ‘potential of hydrogen’) can be used.
Here, the pH scale is the most common and trusted procedure to measure how acidic or basic a substance is.
A pH scale measure can differ from $0 t o 14$ , where $14$ is the most basic, and $0$ is the most acidic a substance can be.
$p H = - l o g [H^{+}]$

For neutral solution, $[H^{+}] = 1 0^{- 7}$ , $p H = 7$ For acidic solution, $[H^{+}] > 1 0^{- 7}, p H < 7$ For basic solution, $[H^{+}] < 1 0^{- 7}, p H > 7$

Note : $pO H = - l o g [O H^{-}]$

$p H + pO H = 14$

-The other way to check if a substance is acidic or basic is by using a litmus paper. There exist two types of litmus paper available, used to identify the acids and bases. They are the red litmus paper and the blue litmus paper. The blue litmus paper changes red under acidic conditions, whereas the red litmus paper turns blue under alkaline or basic conditions.