According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions [`H^+`] and bases are substances that produce hydroxyl ions [`OH^(-)`].

`text(Arrhenius Acids :)` The ionization of an acid `HA` in its aqueous solution can be represented by the following equation

`HA (aq) ⇋ H^(+)(aq) + A^(-)(aq)`

or

`HA(aq) + H_2O(l) ⇋ H_3O^(+) (aq) +A^(-) (aq)`

`text(Important Points)`

- If an acid releases only one `H^(+)` ion per molecule, it is known as monobasic/monoprotic acid (`HA`). For eg. `HCl`, `HBr`, `HI`, `CH_3COOH`, `HNO_3` etc.

- If an acid releases two `H^+` ions per molecule, it is known as dibasic/diprotic (`H_2A`). For eg. `H_2SO_4`, `H_2SO_3`, `H_3PO_3`, `H_2C_2O_4`, `H_2S` etc.

- If an acid releases three `H^+` ions per molecule, it is known as tribasic/triprotic (`H_3A`). For eg. `H_3PO_4`, `H_3AsO_4` etc.

`text(Arrhenius Bases :)` The ionization of a base `BOH ` in its aqueous solution can be represented by the following equation

`BOH (aq) ⇋ B^(+)(aq) + OH^(-)(aq)`

`text(Important Points)`

- If a base releases only one `OH^-` ion per molecule, it is known as monoacidic base `(BOH)`. For eg. `NaOH`, `KOH`, `RbOH`, `CsOH`, `NH_4OH` etc.

- If a base releases two `OH^-` ions per molecule, it is known as diacidic base `[B(OH)_2]`. For eg. `Mg(OH)_2`, `Ca(OH)_2`, `Zn(OH)_2` etc.

- If a base releases three `OH^-` ions per molecule, it is known as triacidic base `[B(OH)_3]`. For eg. `Al(OH)_3`, `Fe(OH)_3` etc.

According to this concept, `HCl` is regarded as an acid only when dissolved in `H_2O` and not in some other solvent such as `C_6H_6` or when it exists in the gaseous form.

According to Bronsted-Lowry theory, an acid is a substance that is capable of donating a hydrogen ion `H^(+)` and bases are substances capable of accepting a hydrogen ion, `H^(+)`. In short, acids are proton donors and bases are proton acceptors.

For eg., the dissolution of ammonia in water can be represented as See fig.1.

In this reaction, water molecule acts as proton donor and ammonia molecule acts as proton acceptor and are thus, called Lowry-Bronsted acid and base, respectively. For the backward reaction, `NH_4^(+)` donates `H^+`, hence it is an acid; `OH^-` accepts `H^+`, hence it is base.

Examples of Bronsted - Lowry Acids & Bases- See fig.2.

`text(Conjugate Acid-Base Pair Concept)` -

An acid-base pair that differs only by one proton is called a conjugate acid-base pair

`text(Consider a reaction)` See fig.3.

In this reaction `HCl` donates a proton to `H_2O` and is, therefore an acid. Water, on the other hand, accepts a proton from `HCl`, and is, therefore, a base. In the reverse reaction, the `H_3O^(+)` ion donates a proton to `Cl^(-)` ion, hence `H_3O^(+)` ion is an acid. `Cl^(-)` ion is a base.

`HCl underset(+H^(+)) overset(-H^(+))⇋ Cl^(-)` and `H_3O^+ underset(+H^(+))overset(-H^(+)) ⇋ H_2O`. Acid base pairs such as which can be formed from each other mutually by the gain or loss of a proton are called conjugate acid - base pairs.

If in the above reaction, the acid `HCl` is an acid and `Cl^-` is its conjugate base.

Similarly, `H_2O` is a base and `H_3O^(+)` is its conjugate acid.

`text(Note-)`

- To get conjugate acid of a given species add `H^(+)` to it. e.g. conjugate acid of `N_2H_4` is `N_2H_5^(+)`.

- To get conjugate base of any species subtract `H^(+)` from it. e.g. Conjugate base of `NH_3` is `NH_2^-`.

- Stronger a Bronsted acid is, weaker is its conjugate base.

- Stronger a Bronsted base is, weaker is its conjugate acid.

More examples -

`text(Acid Conjugate base)`

(i) `HCl quad quad Cl^(-)`

(ii) `H_2SO_4 quad quad HSO_4^(-)`

(iii) `HSO_ 4^(-) quad quad SO_4^(2-)`

(iv) `H_2O quad quad OH^(-)`

`text(Base Conjugate acid)`

(i) `NH_3 quad quad NH_4^(+)`

(ii) `H_2O quad quad H_3O^(+)`

(iii) `RNH_2 quad quad RNH_3^(+)`

`text(Amphiprotic substances)`

Substance that act as an acid as well as a base is called as amphiprotic.

Water can act as an acid in the presence of bases stronger than itself such as `NH_3`, amine, `C_2H_5O^(-)`, `OH^(-)` and `CO_3^(2-)` ions. Water can act as a base in the presence of acids stronger than itself such as `HClO_4`, `HCl`, `CH_3COOH` and phenol.

In fact the amphiprotic nature of `H_2O` is well illustrated in the extremely slight dissociation or self-ionisation :


`oversettext(Weaker acid)(H_2O) +oversettext(Weaker base)(H_2O) ⇋ oversettext(Stronger acid)(H_3O^(+)) + oversettext(Stronger base)(OH^(-))` `(k_w = 1.0 xx 10^(-14))`

According to this theory an acid is any molecule or ion, which can accept an electron pair with the formation of a coordinate bond. For example, in `BF_3` the boron atom can accept a pair of electrons; so `BF_3` is a Lewis acid. A base must therefore be any molecule or ion, which has a lone pair of electrons, which it can donate. For example, ammonia molecule has a lone pair of electrons; so it is a Lewis base. The reaction between a Lewis base and a Lewis acid is just the formation of a coordinate bond between them. See fig.1.

`text(Other examples of Lewis acid-base neutralization)`

`text( Base Acid Products)`

`H_2O + HCl ⇋ H_2O + HCl text(or) H_3O^(+) + Cl^(-)`

`CaO + CO_2 ⇋ CaCO_3 text(or) Ca^(2+) + CO_3^(2-)`

`NH_3 + H_2O ⇋ H_3N HOH text(or) NH_4^(+)- + OH^(-)`

`text(Classification of Lewis Acids :)`

Any Lewis acid must contain at least one empty orbital in the valence shell of one of its atoms to accept an electron pair from a Lewis base. Lewis-acids may be classified as :

i) Molecules containing a central atom with an incomplete octet. Typical examples of this class of acids are electron deficient molecules such as alkyls and halides of `Be`, `B` and `Al`. Some reactions of this type of Lewis acid with Lewis bases are shown below : See fig.2.

Lewis Acid + Lewis base `->` Adducts


ii) `text(Molecules containing a central atom with vacant d-orbitals.)` The central atom of the halides such as `SiX_4, GeX_4, TiCl_4, PCl_3, PF_3, SF_4, SeF_4, TeCl_4` etc. have vacant d-orbitals. These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-orbital and can thus form adducts with a number of halide ions and organic bases. These substances are, therefore, Lewis acids. These halides are vigorously hydrolyzed by `H_2O` to form an oxy acid or oxide of the central atom and the appropriate HX. The hydrolytic reactions take place presumably through the intermediate formation of unstable adducts with `H_2O.` For example

`HF +XeF_7 -> [XeF_5][HF_3]^(-)`

`SbF_5 +2 : underset(. .)overset(. .)F : -> SbF_5^(2-)`

iii) `text(Simple cations )`. Theoretically all simple cations are potential lewis acids. Reactions of some cations as lewis acids with lewis Bases are shown below. It will be seen that these reactions are identical with those which produce Werner complexes. For example, See fig.3.

Ammonation: `Ag^(+) + 2(:NH_3)-> [NH_3-Ag-NH_3]^(+)`

The lewis acid strength or coordinating ability of the simple cations which, according to lewis, are lewis acids, increases with (a) an increase in the positive charge carried by the cation (b) an increase in the nuclear charge for atoms in any period of the periodic
table. (C) a decrease in ionic radius. (d) a decrease in the number of shielding electron shells.

iv) `text(Molecules having multiple bonds between atoms of dissimilar electro-negativity.)` Typical examples of molecules falling in this class of lewis acids are `CO_2, SO_2` and `SO_3`. In these compounds the oxygen atoms are more electronegative than S- or C- atom. As a result, the electron density of p-electrons is displaced away from carbon or sulphur atoms which are less electronegative than oxygen, towards the 0 -atom. The C- or S-atom thus becomes electron deficient and is, therefore, able to accept an electron pair from a Lewis base such as `OH^(-)` ions to from dative bond. See fig.4.