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Definition of Acid and Base

Chemistry Formulas

Definition of Acid and Base

There are several theories to classify acids and bases, which are more-or-less different definitions of what we choose to call an acid or a base. Since it is only a matter of definition, no theory is more right or wrong than any other, and we use the most convenient theory for a particular chemical situation. Of all such theories, lets us take the three important ones.

Arrhenius Concept - The Water Ion System 

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

Arrhenius Acids –

The ionization of an acid HA in its aqueous solution can be represented by the following equation.For More Chemistry Formulas just check out main pahe of Chemsitry Formulas. 

HA (aq)  H+ (aq) + A- (aq)


HA(aq) + H2O(l)  H3O+ (aq) + A- (aq)

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, CH3COOH, HNO3 etc.
  • If an acid releases two H+ ions per molecule, it is known as dibasic/diprotic (H2A). For eg. H2SO4, H2SO3, H3PO3 ,H2C2O4 , H2S etc.
  • If an acid releases three H+ ions per molecule, it is known as tribasic/triprotic (H3A). For eg. H3PO4, H3AsO4 etc.

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)

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, NH4OH 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 H2O and not in some other solvent such as C6H6 or when it exists in the gaseous form

Brönsted - Lowry Concept  (The Proton - donor - Acceptor Concept)

According to Brönsted-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

bronsted lowry concept

Conjugate Acid - Base Pair Concept

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

Consider a reaction

conjugate acid base pair concept

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, H2O is a base and H3O+ is its conjugate acid.

Note –

·    To get conjugate acid of a given species add H+ to it. e.g. conjugate acid of N2H4 is N2H5+.

·    To get conjugate base of any species subtract H+ from it. e.g. Conjugate base of NH3 is NH2-.

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

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

More examples –

     Acid                                                 Conjugate base

(i)         HCl                                                            Cl-

(ii)        H2SO4                                             HSO4-

(iii)       HSO4-                                                      SO42-

(iv)       H2O                                                OH-

Base                                                Conjugate acid

(1)          NH3                                               NH4+

(2)          H2O                                               H3O+

(3)          RNH2                                             RNH3+


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 NH3, amine, C2H5O, OH and CO32– ions.  Water can act as a base in the presence of acids stronger than itself such as HClO4, HCl, CH3COOH and phenol.

In fact the amphiprotic nature of H2O is well illustrated in the extremely slight dissociation or self-ionisation:

amphirprotic substance

The Lewis Concept (Electron Donor - Acceptor System)

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 BF3 the boron atom can accept a pair of electrons; so BF3 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.

lewis concept

Other examples of Lewis acid-base neutralization

lewis acid-base neutralization

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 form 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:

  classification of lewis acids

ii) Molecules containing a central atom with vacant d-orbitals. The central atom of the halides such as SiX4, GeX4, TiCl4, PCl3, PF3, SF4, SeF4, TeCl4 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 H2O 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 H2O. For example 

lewis acid

iii) 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,

simple cations

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) Molecules having multiple bonds between atoms of dissimilar electro-negativity. Typical examples of molecules falling in this class of Lewis acids are CO2, SO2 and SO3. 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 O-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.

lewis acid



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