JEE Main & Advanced Chemistry Equilibrium / साम्यावस्था Acid And Bases

Acid And Bases

Category : JEE Main & Advanced

(1) Arrhenius concept : According to Arrhenius concept all substances which give H+ ions when dissolved in water are called acids while those which ionise in water to furnish OH ions are called bases.

\[\underset{(Acid)}{\mathop{HCl}}\,\]        \[\underset{(aq.)}{\mathop{{{H}^{+}}}}\,+\underset{(aq.)}{\mathop{C{{l}^{-}}}}\,\] ;   \[\underset{(Base)}{\mathop{NaOH}}\,\]         \[\underset{(aq.)}{\mathop{N{{a}^{+}}}}\,+\underset{(aq)}{\mathop{O{{H}^{-}}}}\,\]

Some acids and bases ionise almost completely in solutions and are called strong acids and bases. Others are dissociated to a limited extent in solutions and are termed weak acids and bases. \[HCl,\ HN{{O}_{3}},\ {{H}_{2}}S{{O}_{4}},\,HCl{{O}_{4}}\], etc., are examples of strong acids and \[NaOH,\ KOH,\ {{(C{{H}_{3}})}_{4}}NOH\] are strong bases. Every hydrogen compound cannot be regarded as an acid, e.g., \[C{{H}_{4}}\] is not an acid. Similarly, \[C{{H}_{3}}OH,\ {{C}_{2}}{{H}_{5}}OH\], etc., have \[OH\] groups but they are not bases.

(i) Utility of Arrhenius concept : The Arrhenius concept of acids and bases was able to explain a number of phenomenon like neutralization, salt hydrolysis, strength of acids and bases etc.

(ii) Limitations of Arrhenius concept

(a) 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.

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

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

     (d) It cannot explain the acidic character of certain salts such as \[AlC{{l}_{3}}\] in aqueous solution.


(2) Bronsted–Lowry concept : According to this concept,

“An acid is defined as a substance which has the tendency to give a proton (H+) and a base is defined as a substance which has a tendency to accept a proton. In other words, an acid is a proton donor whereas a base is a proton acceptor.”

\[\underset{\text{Acid}}{\mathop{HC{{l}_{{}}}}}\,\,+\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\] ? \[{{H}_{3}}{{O}^{+}}+C{{l}^{-}}\]                                                                                                                                                               …..(i)

\[\underset{\text{Acid}}{\mathop{C{{H}_{3}}COOH}}\,\,+\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\] ? \[{{H}_{3}}{{O}^{+}}+C{{H}_{3}}CO{{O}^{-}}\]                                                                                                                     …..(ii)

(i) \[HCl\] and \[C{{H}_{3}}COOH\] are acids because they donate a proton to \[{{H}_{2}}O\].(ii) \[N{{H}_{3}}\] and \[CO_{3}^{2-}\] are bases because they accept a proton from water.

In reaction (i), in the reverse process, H3O+ can give a proton and hence is an acid while Cl can accept the proton and hence is a base. Thus there are two acid-base pairs in reaction (i). These are HClCl and H3O+H2O. These acid-base pairs are called conjugate acid-base pairs.

Conjugate acid ? Conjugate base \[+{{H}^{+}}\]

Conjugate base of a strong acid is a weak base and vice a versa. Weak acid has a strong conjugate base and vice a versa.

Levelling effect and classification of solvents : In acid-base strength series, all acids above H3O+ in aqueous solution fall to the strength of H3O+. Similarly the basic strength of bases above OH fall to the strength of OHin aqueous solution. This is known as levelling effect. Levelling effect of water is due to its high dielectric constant and strong proton accepting tendency.

On the basis of proton interaction, solvents are of 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.

HCl acts as acid in H2O, stronger acid in NH3, weak acid in CH3COOH, neutral in C6H6 and a weak base in HF.

\[\underset{Base}{\mathop{HCl}}\,\,+\,\underset{Acid}{\mathop{HF}}\,\,\,\,\to \,\,\,\underset{Acid}{\mathop{{{H}_{2}}C{{l}^{+}}}}\,\,+\,\underset{Base}{\mathop{F_{{}}^{-}}}\,\]

Utility of Bronsted – Lowry concept

(i) Bronsted – Lowry concept is not limited to molecules but includes even the ionic species to act as acids or bases.

(ii) It can explain the basic character of the substances like \[N{{a}_{2}}C{{O}_{3}},\ N{{H}_{3}}\] etc.

(iii) It can explain the acid-base reactions in the non-aqueous medium or even in the absence of a solvent               (e.g., between HCl and NH2).

Limitations of Bronsted lowry concept

(i) The protonic definition cannot be used to explain the reactions occuring in non-protonic solvents such as COCl2, SO2, N2O4, etc.

(ii) It cannot explain the reactions between acidic oxides like \[C{{O}_{2}},\ S{{O}_{2}},S{{O}_{3}}\] etc and the basic oxides like \[CaO,\ BaO,\ MgO\] etc which take place even in the absence of the solvent e.g.,

\[CaO+S{{O}_{3}}\to CaS{{O}_{4}}\]

There is no proton transfer in the above example.

(iii) Substances like BF3, AlCl3 etc, do not have any hydrogen and hence cannot give a proton but are known to behave as acids.

 Conjugate acid-base pairs

 (3) Lewis concept : This concept was proposed by G.N. Lewis, in 1939. According to this concept, “a base is defined as a substance which can furnish a pair of electrons to form a coordinate bond whereas an acid is a substance which can accept a pair of electrons.” The acid is also known as electron pair acceptor or electrophile while the base is electron pair donor or nucleophile.

A simple example of an acid-base is the reaction of a proton with hydroxyl ion,      \[\underset{Acid}{\mathop{{{H}^{+}}}}\,+\underset{Base}{\mathop{O{{H}^{-}}}}\,\to HOH\]

Lewis concept is more general than the Bronsted Lowry concept. All Bronsted bases are also Lewis bases but all Bronsted acids are not Lewis acids. [e.g., \[HCl,\,{{H}_{2}}S{{O}_{4}}\]as they are not capable of accepting a pair of electrons]

(i) Types of Lewis acids : According to Lewis concept, the following species can act as Lewis acids.

(a) Molecules in which the central atom has incomplete octet \[B{{F}_{3}},\ BC{{l}_{3}},\ AlC{{l}_{3}},BeC{{l}_{2}}\], etc.

(b) All cations are expected to act as Lewis acids since they are deficient in electrons.

(c) Molecules in which the central atom has empty              d-orbitals. e.g., \[Si{{F}_{4}},\,SnC{{l}_{4}},P{{F}_{5}}\]etc.

(d) Molecules having a multiple bond between atoms of dissimilar electronegativity e.g., \[C{{O}_{2}},S{{O}_{2}}\].  

(ii) Types of Lewis bases : The following species can act as Lewis bases.

(a) Neutral species having at least one lone pair of electrons \[:\,N{{H}_{3}},\,-\overset{.\,\,.}{\mathop{N}}\,{{H}_{2}},\,R-\underset{.\,\,.}{\mathop{\overset{.\,\,.}{\mathop{O}}\,}}\,-H\]

(b) Negatively charged species or anions

(iii) Hard and Soft principle of acids and bases : Lewis acids and bases are classified as hard and soft acids and bases. Hardness is defined as the property of retaining valence electrons very strongly. Thus a hard acid is that in which electron-accepting atom is small, has a high positive charge and has no electron which can be easily polarised or removed e.g., \[L{{i}^{+}},N{{a}^{+}},B{{e}^{2+}},M{{g}^{+2}},A{{l}^{+3}}\]  \[B{{F}_{3}},S{{O}_{3}}\] etc.. On the contrary, a soft acid is that in which the acceptor atom is large, carries a low positive charge or it has electrons in orbitals which are easily polarised or distorted         e.g., \[P{{b}^{+2}},C{{d}^{+2}},P{{t}^{+2}},Hg_{{}}^{+2},R{{o}^{+}},R{{s}^{+}},{{I}_{2}}\] etc..

A Lewis base which holds its electrons strongly is called hard base, e.g., \[O{{H}^{-}},{{F}^{-}},{{H}_{2}}O,N{{H}_{3}},C{{H}_{3}}OC{{H}_{3}}\], etc. on the other hand, a Lewis base in which the position of electrons is easily polarised or removed is called a soft base e.g., \[{{I}^{-}},CO,\ C{{H}_{3}}{{S}^{-}},{{(C{{H}_{3}})}_{3}}P\], etc.

In general, hard acids prefer to bind to hard bases and soft acids prefer to bind to soft bases. The bonding between hard acids and hard bases is chiefly ionic and that between soft bases and soft acids is mainly covalent.

(iv) Utility of Lewis concept : Lewis concept is the most general of all the concepts and can explain the acidic and basic nature of all those substances which could not be explained by the earlier concepts. Similarly, it can explain even those acid-base reactions which could not be explained by the other concepts.

(v) Limitations of lewis concept : It does not explain behaviour of well known protonic acids, as \[HCl,\,{{H}_{2}}S{{O}_{4}}\] etc, as which do not form coordinate bonds with bases.

It does not explain relative strengths of acids and bases. Many lewis acids do not posses catalytic property..

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