Acids and Bases:

The earliest criteria for the characterization of acids and bases were the experimentally observed properties of aqueous solutions. An acid* was defined as a substance whose water solution tastes sour, turns blue litmus red, neutralizes bases and so on. A substance was a base if its aqueous solution tasted bitter, turns red litmus blue, neutralizes acids and so on. Faraday termed acids, bases and salts as electrolytes and Liebig proposed that acids are compounds containing hydrogen that can be replaced by metals.

Different concepts have been put forth by different investigators to characterize acids and bases but the following are the three important modern concepts of acids and bases:

(1)  Arrhenius concept

 According to Arrhenius concept all substances which give H+ ions when dissolved in water are called acids while those which ionize in water to furnish OH- ions are called bases.

            HA ↔  H+ + A- (Acid)

            BOH ↔ B+ + OH- (Base)


Thus, HCl is an acid because it gives H+ ions in water. similarly, NaOH is a base as it yields OH- ions in water.

            HCl ↔ H+ + Cl-

            NaOH ↔  Na+ + OH-


Some acids and bases ionize completely in solutions and are called string acids and bases. Others are dissociated to a limited extent in solutions and are termed weak acids and bases. HCl, HNO3, H2SO4, HCIO4, etc., are examples of strong acids and NaOH, KOH, (CH3)4NOH are strong bases. Every hydrogen compound cannot be regarded as an acid, e.g., CH4 is not an acid. Similarly, CH3OH, C2H5OH, etc., have OH groups but they are not bases.

Actually free H+ ions do not exist in water. they combine with solvent molecules, i.e., have strong tendency to get hydrated.

                    HX + H2O ↔ H3O+ + X-

                                                   (Hydronium ion)


The proton in aqueous solution is generally represented as H+ (aq). It is now known that almost all the ion are hydrated to more or less extent and it is customary to put (aq) after each ion.


The oxides of many non-metals react with water to form acids and are called acidic oxides or acid anhydrides.

            CO2 + H2O  H2CO3 ↔ 2H+(aq) +  (aq)

            N2O5 + H2O  2HNO3 ↔ 2H+(aq) +  (aq)


Many oxides of metals dissolve in water to form hydroxides. Such oxides are termed basic oxides.

Na2O + H2O →  2NaOH  ↔  2Na+(aq) + 2OH- (aq)


The substance like NH3 and N2H4 act as bases as they react with water to produce OH- ions.

NH3 + H2O →  NH4OH  ↔  NH+4 (aq) + OH- (aq)


The reaction between an acid and a base is termed neutralization. According to Arrhenius concept, the neutralization in aqueous solution involves the reaction between H+ and OH- ions or hydronium and OH-. This can be represented as

        H3O+ + OH- ↔  2H2O



(i)  For the acidic or basic properties, the presence of water is absolutely necessary. Dry HCl shall not act as an acid. HCl is regarded as an acid only when dissolved in water and not in any other solvent.

(ii)  The concept does not explain acidic and basic character of substances in non-aqueous solvents.

(iii) The neutralization process is limited to those reactions which can occur in aqueous solutions only, although reactions involving salt formation do occur in the absence of solvent.

(iv) It cannot explain the acidic character of certain salts such as AlCl3 in aqueous solution.

(v) An artificial explanation is required to explain the basic nature of NH3 and metallic oxides and acidic nature of non-metal oxides.


(2)  Bronsted-Lowry concept - The proton-donor-acceptor concept:

In 1923, Bronsted and Lowry independently proposed a broader concept of acids and bases. According to Bronsted-Lowry concept an acid is a substance (molecule or ion) that can donate proton, i.e., a hydrogen ion, H+, to some other substance and a base is a substance that can accept a proton from an acid. More simply, an acid is a proton donor (protogenic) and a base is a proton acceptor (protophilic).

Consider the reaction,

                    HCl + H2O ↔  H3O + Cl-

In this reaction HCl acts as an acid because it donates a proton to the water molecule. Water, on the other hand, behaves as a base by accepting a proton from the acid.

    The dissolution of ammonia in water may be represented as

                    NH3 + H2O ↔ NH4+ + OH-

In this, reaction, H2O acts as an acid it donated a proton to NH3 molecule and NH3molecule behaves as a base as it accepts a proton.

When an acid loses a proton, the residual part of it has a tendency to regain a proton. Therefore, it behaves as a base.

                Acid ↔ H+ + Base

The acid and base which differ by a proton are known to form a conjugate pair. Consider the following reaction.

 CH3COOH + H2O ↔  H3O+ + CH3COO-

It involves two conjugate pairs. The acid-base pairs are:


Such pairs of substances which can be formed form one another by loss or gain of a proton are known as conjugate acid-base pairs.

If in the above reaction, the acid CH3COOH is labelled acid1 and its conjugate base, CH3COO- as base1. H2O is labelled as base2 and its conjugate acid H3O+ as acid2, the reaction can be written as:

                Acid1 + Base2 ↔ Base1 + Acid2


Thus, any acid-base reaction involves two conjugate pairs, i.e., when an acid reacts with a base, another acid and base are formed. Some more examples are given below:


Thus, every acid has its conjugate base and every base has its conjugate acid. It is further observed that strong acids have weak conjugate bases while weak acids have strong conjugate bases.

HCl                   Cl-               CH3COOH        CH3COO-

Strong acid       Weak base   Weak acid        Strong base


There are certain molecules which have dual character of an acid and a base. These are called amphiprotic or atmospheric.

Examples are NH3, H2O, CH3COOH, etc.


The strength of an acid depends upon its tendency to lose its proton and the strength of the base depends upon its tendency to gain the proton.


   Acid-Base chart containing some common conjugate acid-base pairs



Conjugate base


(Perchloric acid)


(Perchlorate ion)


(Sulphuric acid)


(Hydrogen sulphate ion)


(Hydrogen chloride)


(Chloride ion)


(Nitric acid)


(Nitrate ion)


(Hydronium ion)




(Hydrogen sulphate ion)


(Sulphide ion)


(Ortho phosphoric acid)


(Dihydrogen phosphate ion)


(Acetic acid)


(Acetate ion)


(Carbonic acid)


(Hydrogen carbonate ion)


(Hydrogen sulphide)


(Hydorsulphide ion)


(Ammonium ion)




(Hydrogen cyanide)


(Cyanide ion)


(Ethyl alcohol)


(Phenoxide ion)




(Hydroxide ion)




(Ethoxide ion)




(Amide ion)




(Methide ion)


In acid-base strength series, all acids above H30+ in aqueous solution fall to the strength of H30+. Similarly the basic strength of bases below OH" fall to the strength of OH" in aqueous solution. This is known as levelling effect.

The strength of an acid also depends upon the solvent. The acids HCIO4, H2S04, HCl and HN03 which have nearly the same strength in water will be in the order of HC104> H2S0> HCl > HN03 in acetic acid, since the proton accepting ten­dency of acetic acid is much weaker than water. So the real strength of acids can be judged by solvents. On the basis of proton interaction, solvents can be classified into four types:


(i)  Protophilic solvents:  Solvents which have greater tendency to accept protons, i.e., water, alcohol, liquid ammonia, etc.

(ii) Protogenic solvents:  Solvents which have the tendency to produce protons, i.e., water, liquid hydrogen chloride, glacial acetic acid, etc.

(iii) Amphiprotic solvents:  Solvents which act both as protophilic or protogenic, e.g., water, ammonia, ethyl alcohol, etc.

(iv) Aprotic solvents:  Solvents which neither donate nor accept protons, e.g., benzene, carbon tetrachloride, carbon disulphide, etc.


HCI acts as acid in H20, stronger acid in NH3, weak acid in CH3COOH, neutral in C6H6and a weak base in HF.

HCI   +   HF     →    H2C1+  +   F-

Base      Acid          Acid             Base

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