Relative Strengths of Acids & Bases

We have already discussed several definitions of acids and bases. It is interesting to study the general trends in acid or base strength among such compounds.

The protonic acids may be broadly classified as

(i) hydroacids and

(ii) oxo acids.

In hydroacids, the acidic hydrogen atom is directly linked to some second element. Binary hydrides like HF, HCl, H2S are such examples. In oxo acids, the acidic hydrogen is linked to some central atom via an oxygen atom e.g. for H2SO4, HNO3 or HClO4.

According to Bronsted concept, a stronger acid has a stronger tendency to donate a proton and a strong base has a strong tendency to accept a proton. At least two general methods are generally used for the comparison of relative acidity of given acids.

(i) The first of these consists of making a comparison of proton-donating tendencies of different acids towards the same base. For moderately strong acids, H2O is generally used as the base.

Suppose we have to compare the acidic strengths of CH3COOH and HCN. Experimentally it has been observed that the ionisation or acidity constant, Ka for CH3COOH and HCN at 25°C is 1.8 × 10-5 and 4.0 × 10-10 respectively.

CH3COOH + H2O <–> H3O+ + CH3COO (Ka = 1.8 × 10-5)

HCN + H2O <–> H3O+ + CN (Ka = 4.0 × 10-10)

CH3COOH is, therefore, a stronger acid than HCN and CN ion is a stronger base than CH3COO ion.

(ii) The second method is the competitive protolysis method. In this method one acid is added to the conjugate base of another and the equilibrium concentration are determined experimentally. For example, when NaOC2H5 is added to H2O, it is
experimentally seen that ion, which is the conjugate base of C2H5OH reacts fairly completely with H2O to form C2H5OH and OH ion.

Ethoxide ion, C2H5O is, therefore, a stronger base than OH and H2O is a stronger acid than C2H5OH.

Similarly when HS is added to NH3, it has been found experimentally that NH4+ and S2 ions are present in the reaction mixture. This shows that NH3 is a stronger base in comparison to HS

Periodic variations of acidic and basic properties

The discussion of this topic is made under the following heads.

(a) Hydracids of the elements of the same period. We can consider the hydracids of the elements of 2nd period viz. CH4, NH3, H2O and HF. These hydrides become increasingly acidic as we move from CH4 to HF.

Thus CH4 has negligible acidic properties, but NH3 donates a proton (H+) to strong bases to form , H2O loses a proton even more readily and HF is a fairly strongly acid. The increase in the acidic properties of these hydrides is due to the fact that as we move from CH4 to HF, the stability of their conjugate base viz , , OH and F increases in the order:

CH3 < NH3 < OH < F

the increase in acidic properties is supported by the successive increase in the dissociation constant values of these hydrides as shown.

CH4(=10-58) < NH3(=10-35) < H2O(= 10-14) < HF (= 10-4)

(b) Hydracids of the elements of the same group. The following examples make this point clear:

(i) Hydracids of VA groups elements (NH3,PH3, AsH3, SbH3,BiH3). All these hydrides show basic character. With the increase in size and decrease in electronegativity from N to Bi, there is a decrease in electron density in sp3 hybrid orbital and thus electron donor capacity (i.e.basic character) decreases.

(ii) Hydracids of VI A group elements (H2O, H2S, H2Se, H2Te). Aqueous solution of the hydrides of this group behave as weak diprotic acid and ionise as:

H2R <–> H+ + HR

HR <–> H+ + R2

The strength of the hydrides as acids increases in the order:

H2O < H2S < H2Se < H2Te.

This order is supported by the successive increase of their dissociation constants as shown.

H2O (1.0 × 10-14) < H2S (1.1× 10-7) < H2Se(2 × 10-4) < H2Te(2.3 × 10-3)

The increasing acidic character reflects decreasing trend in the electron donor ability of OH, HS,HSeor HTe ions. The increasing acidic character is explained by saying that as the charge density on the conjugate base decreases from OH to HTe , the proton is less tightly held in higher members and, therefore, acidic character increase.

(iii) Hydracids of VIIA group elements (HF, HCl, HBr, HI). The aqueous solutions of these hydrides show acidic character which increases in the order
HF < HCl < HBr < HI.

This order is explained by saying that as we pass from HF to HI, there is a gradual decrease in the bond energies of H – X bonds (H – F = 135 kcal/mole, H-Cl = 103, H – Br = 88, H – I = 71).

This decreasing order of bond energies increase the tendency to HX molecule to split up into H+ and Xions in aqueous solution and thus the acidic character increases from HF to HI.

Oxyacids

(c) Oxyacids.
(i) The acidic character of oxyacids of the same element which is in different oxidation states increases with the increase of its oxidation state. The following series follow this rule (called oxidation number or oxidation state rule).

(a) HCl+O < HCl3+O2 < HCl5+O3 < HCl7+O4

b) H2S4+O3 < H2S6+O4 (c) HN3+O2 < HN5+O3.

Explanation: With reference to the oxyacids of halogens explanation of the oxidation rule can be given as follows. It is well known that the stronger the acid, the weaker will be its conjugate base and vice versa.
Now the conjugate bases of the acids are :
ClO, ClO2 , ClO3 , ClO4 respectively the oxyanions in which the central atom (i.e. chlorine atom) has larger oxidation number, has the larger number of lone oxygen atoms for participation in extension of the π bond. Thereby the charge on the ion is delocalised which greatly stabilises the ion and thus deceases its tendency to accept a proton i.e., causes the ions to be a very weak base with the result that the strength of the acid increases.

When the oxidation state rule as given above is applied to the oxyacids of phosphorus viz. H3P+O2 < H3P+3O3 < H3P+5O4, but the experimental observation suggest the reverse order viz H3PO2 ≥ H3PO3 > H3PO4.

Explanation: The experimental order can be explained when we consider the structures of these acids as given below. In these the number of protonated and unprotonated oxygen atoms have also been indicated. The oxygen atom attached with a proton is called protonated oxygen while that attached directly with phosphorus (central atom) is known as unprotonated oxygen.

The proton attached to any oxygen atom has a far greater chance of dissociation than that linked directly with phosphorus atoms (which is the central atom). Thus in this series, since the number of protonated oxygen atoms and consequently the number of dissociable protons increases from one in H3PO2 to three in H3PO4, the acidity of these acids decrease in the order : H3PO2 ≥ H3PO3 > H3PO4

(ii) The acidic character of the oxyacids of different elements which are in the oxidation state decrease with the increase in atomic number of the central atom. The following series follow this rule

a) HOCl+ > HOBr+ > HOI+

b) HCl7+ O4 > HI7+O4

c) H2S4+O3 > H2Se4+O3

Explanation: As the atomic number of the central atom increase, its electronegativity decreases and its size increase. As a result of this the tendency of the acid to lose a proton to water decrease. This makes the acid a weaker acid.

d) Hydrated metal ions: Under favourable conditions one or more protons may dissociate from the coordinated aquo groups.

Thus hydrated metal ions also develop acidity. The other things being equal, acidity increases with the increase of positive charge and basicity increases with the increase of negative charge. Thus [Fe(H2O)6]3+ ion is a stronger acid than [Fe(H2O)6]2+ ion and [Ni(OH)4]2 is a stronger base than [Ni(OH)4] ion.

Amphiprotonic substances

These substances act as an acid as well as a base: e,g, CH3COOH is acid with water while is a base with HF.

Urea is also anphiprotonic as it is an acid with NH3 while a base with sulphuric acid

Similarly 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 amphiprotonic nature of H2O is well illustrated in the extremely slight dissociation or self-ionisation:

(Kw = 1.0 × 10-14)

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