There are several so-called theories of acids and bases, but they are not really theories but merely 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. So before we talk of strength of acids and bases, we need to know several theories.
Modern Concepts of Acids and Bases :
Following are the important modern concept of acids and bases:
Arrhenius Concept – The Water Ion System
According to this concept, an acid is any hydrogen containing compound which gives H+ ions in aqueous solution and a base which gives OH– ions in aqueous solution. The HCl is an acid and NaOH is a base and the neutralisation process can be represented by a reaction involving the combination of H+ and OH– ions to form H2O.
NaOH ⇔ Na+ + OH–
H+ + OH– → H2O
(i) Since the reaction representing neutralisation process involves the combination of H+ and OH– ions, the approximately constant molar heat of neutralisation would be expected. Thus the constant heat of neutralisation of a strong acid by a strong base is readily understandable in terms of this concept.
(ii) This concept has offered a means of correlating catalytic behaviour with the concentration of the H+ion.
(i) The acid or base property of a substance was not inherent in it, but was dependent on the presence of water. 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.
(ii) The neutralisation process is limited to those reactions which can occur in aqueous solution only, although reactions involving salt formation do occur in many other solvents and even in the absence of solvents.
(iii) It cannot explain the acidic character of certain salts such as AlCl3 in aqueous solution.
Bronsted – Lowry Theory The Proton – donor – Acceptor System
Bronsted and Lowry in 1923 independently proposed a more general definition of acids and bases. According to them, an acid is defined as any hydrogen containing material (a molecule or a cation or an anion) that can release a proton (H+) to any other substance, whereas a base is any substance (a molecule or a cation or an anion) that can accept a proton from any other substance. In short, an acid is a proton -donor and a base is a proton – acceptor.
Conjugate Acid – Base Pairs :
Consider a reaction
Acid1 Base2 Acid2 Base1
H2O + NH3 ⇔ H3O+ OH–
In this reaction HCl donates a proton to H2O and is, therefore an acid. Water, on the other hand, accepts a proton from HCl, and is, therefore, a base. In the reverse reaction which at equilibrium proceeds at the same rate as the forward reaction, the H3O+ ions donates a proton to Cl– ion, hence H3O+, ion is an acid. Cl– ion, because it accepts a proton from H3O+ ion, is a base. Acid base pairs such as.
the members of which can be formed from each other mutually by the gain or loss of proton are called conjugate acid – base pairs.
If in the above reaction, the acid HCl is labelled Acid1 and its conjugate base viz. Cl– as Base1 and further, if H2O is designated Base2 and its conjugate acid viz. H3O+ as Acid2, the equilibrium can be represented by a general equation.
This is the fundamental equation representing the relationship between an acid and a base on the basis of Bronsted concept. Thus on the basis of this concept Acid1 and Base1 form one conjugate acid-base pair and Acid2 and Base2 form another conjugate acid-base pair
Axioms of the Bronsted concept & position of equilibrium in acid-base reactions
HCl + H2O ⇔ H3O+ + Cl–
In the equilibrium mixture two acid HCl and H3O+ ion are competing to donate protons to a base.
Since HCl wins, it is the stronger acid. Similarly two bases, H2O and Cl– ion, are competing to accept protons.
Here H2O is the stronger base. It will be seen that the stronger acid, HCl, has the weaker conjugate base Cl– ion and the stronger base, H2O, has weaker conjugate acid, H3O+ ion.
The stronger acid and weaker base form one conjugate acid – base pair and weaker acid and stronger base form another conjugate acid base pair. It is quite evident that HClO4 is the strongest acid; its conjugate base ion, is consequently the weakest base. CH4 and H2 are the weakest acids; their conjugate bases, ion and H– ion respectively, are consequently the strongest bases.
As a stronger acid, HCl is highly ionised even in concentrated aqueous solution. At equilibrium, the above reaction proceeds to the right, with most of HCl ionised to form H3O+ and Cl– ions. This fact can be illustrated by using arrows of unequal length to designate the forward and reverse reactions respectively. Thus.
Stronger acid + Stronger Base <=> Weaker acid + Weaker Base
HCl + H2O <=> H3O+ + Cl– ……(1)
The longer arrow indicates that the position of equilibrium lies to the right.
In the ionisation of CH3COOH in H2O, equilibrium is reached when the reaction has proceeded to the right only to slight extent, with only a small fraction of the CH3COOH present in the form of ions.
Weaker acid + Weaker base <=> Stronger acid + Stronger base
CH3COOH + H2O <=> H3O+ + CH3COO– ….. (2)
Here the longer arrow indicates that the position of equilibrium lies to the left.
Evidently H3O+ ion in equilibrium (2) is a stronger acid and CH3COO– ion is a stronger base. It is also evident that the stronger acid H3O+ ion has the weaker conjugate base, H2O and the stronger base, CH3COO– has the weaker conjugate acid, CH3COOH.
We thus see that all the proton transfer reactions (i.e., protolysis reactions) run down hill to form predominantly the weaker acid and the weaker base.
Also Read :
| Relative Strengths of Acids & Bases
Levelling effect , levelling & differentiating solvents
General theory of solvent system
Hard and Soft Acids and Bases
Origin of Acidity & Basicity in Organic Compounds