Concepts of Acids and Bases_Final
- IIT-JEE Syllabus
Various concepts of Acids and Bases, Relative Strengths of Inorganic and Organic Acids and Bases.
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:
3.1 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+
- i) According to this concept, HCl is regarded 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) It cannot account for the acidic and basic character of the materials in non-aqueous solvents, e.g. NH4NO3 in liq. NH3 acts as an acid, though it does not give H+ Similarly many organic materials and NH3, which do not have OH– ions at all, are actually known to show basic character.
iii) 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.
- iv) It cannot explain the acidic character of certain salts such as AlCl3 in aqueous solution.
3.2 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 NH4+ 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.
Two important axioms of the Bronsted concept and position of equilibrium in acid-base reactions:
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 downhill to form predominantly the weaker acid and the weaker base.
Relative Strengths of Acids and Bases
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° is 1.8 × 10–5 and 4.0 × 10–10
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–
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 strong 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:
< < 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,SH3, 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–,HSe– or 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 kcal/mole H – Br = 88, H – I = 71). This decreasing order of bond energies increase the tendency of HX molecule to split up into H+ and X– ions in aqueous solution and thus the acidic character increases from HF to HI.
- 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–, , , 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.
|Nature of the acid:||Monobasic||Diabsic||Tribasic|
|No. of protonated oxygen atoms||1||2||3|
|No. of unprotonated oxygen atoms:||1||1||1|
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.
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 amphiprotic nature of H2O is well illustrated in the extremely slight dissociation or self-ionisation:
+ + (Kw = 1.0 × 10–14)
The levelling effect, levelling and differentiating solvents:
The apparent strength of a protonic acid is dependent on the solvent in which the acid is dissolved. When all the acids in the acid chart which are stronger than H3O+ ion (i.e., the acids above H3O+ acids) are added to H2O, they donate as proton to H2O to H3O+ ion and appear to have equal strength, since all these acids are levelled to the strength of H3O+ ion which is left in solution and is common to all such solutions. This phenomenon viz. The strength of all the acids becomes equal to that of H3O+ ion is called leveling effect of the solvent, water, and here water is called a leveling solvent for all these acids.
In aqueous solution all very strong bases like Na+H–, Na+NH2–, Na+OC2H5– are levelled to the strength of OH– ion, for they react completely with H2O to produce OH– ions. The solvent in which complete proton-transfer occurs are called levelling solvents. In other words, the solvent in which the solute is ~100% ionised, are called levelling solvents. Since HF and HCl both are ~ 100% ionised in liquid NH3 to give ~100% NH4+ ions, these appear to be of equal strength and liq. NH3 acts as a levelling solvent for HF and HCl. In H2O, HF is only partially ionised, whereas HCl and HBr are ~ 100% ionised. Thus H2O is a differentiating solvent for HF, but for HCl and HBr it is a leveling solvent. Several mineral acids are partially ionised in glacial CH3COOH medium because CH3COOH is a poor proton-acceptor but rather a better proton donor. CH3COOH, therefore, acts as a differentiating solvent towards the mineral acids. But, for bases, CH3COOH act as a levelling solvent.
Utility of Bronsted Concept
- i) It defines acids and bases in terms of the substances themselves and not in terms of their ions in aqueous solution. Thus HCl is an acid because of the fact that it can give a H+
HCl H+ + Cl–
- ii) The Bronsted concept recognises that acid-base behaviour is neither restricted to, nor depends on, any particular, solvent
iii) This concept is useful in accounting for the hydrolysis of salt solution. When a salt is dissolved in water, an inbalance in the concentration of the solvent cation (H3O+) and anion (OH–) will result, if the salt cation and anion differ in their proton-donor and proton-acceptor properties towards H2O. This point can be illustrated by considering the aqueous solution of FeCl3. Aqueous hydrated ferric ion, [Fe(H2O)x]3+, exceeds the proton-acceptor ability of Cl– ion and a considerable excess of H3O+ ion is produced in the solution, making FeCl3 acidic.
FeCl3 + xH2O [Fe(H2O)x]3+ + 3Cl–
Aqueous solution of a Na2CO3 is alkaline in character, because the proton acceptor ability of ion exceeds the proton-donor ability of hydrated sodium ion, [Na(H2O)x]+.
2xH2O + Na2CO3 2[Na(H2O)x]+ +
- i) This concept lays excessive emphasis on the proton – transfer. Although it is true that most common acids are protonic in nature, yet there are many which are not.
- ii) There are a number of acid-base reactions in which no proton transfer takes place, e.g. SO2 + SO2 SO2+ +
Acid1 + Base2 Acid2 + Base1
Thus the protonic definition cannot be used to explain the reactions occuring in non-protonic solvents such as COCl2,SO2, N2O4 and BrF3.
Exercise 1 Arrange according to increasing Lewis acid character,
SiF4, SiCl4, SiBr4, Sil4
3.3 General theory of solvent system
The protonic definition of acids and bases given by Bronsted can be extended to the reactions occurring in non-aqueous solvents containing hydrogen such as NH3, N2H4,HF, H2SO4, CH3COOH, HCN and alcohols.
In an attempt to have a more general definition of acids and bases applicable to protonic and non-protonic solvents, several definitions have been proposed. One of these is due to Cady and Elsey (1928) according to whom an acid is solute that, either by direct dissociation or by reaction with the solvent gives the anion characteristic of the solvent and a base is a solute that, either by direct association or by reaction with the solvent, gives the cation characteristic of the solvent. If for example, we consider the solvent H2O, its characteristic cation and anion are H3O+ and OH– respectively as shown below:
Thus all those compounds which can give H3O+ ions in H2O will act as acids and all the compounds which can give OH– ions in H2O will behave as bases.
Similarly in N2O4 as solvent substance such as NOCl which yield NO+ ions are acids and the substances such as NaNO3 which yield NO3– ions are bases.
Evidently this definition of acids applies equally well to protonic and non-protonic solvents.
The auto-ionisation of some protonic and non-protonic solvents is shown below.
Just as with the Arrhenius definition, neutralisation is a reaction between an acid and a base to produce a salt and the solvent. Neutralisation reaction in some non-aqueous solvents are given below:
|In liq. NH3||:||NH4Cl||+||NaNH2||NaCl||+||2NH3|
|In liq. SO2||:||SOCl2||+||[N(CH3)4]2 SO3||2[N(CH3)4]Cl||+||2SO2|
It may be seen from the following reactions that there is a complete analogy between the solvolytic acid and amphoteric behaviours in aqueous solvents.
|Solvolytic behaviour||In liq. NH3:||AlCl3 + NH3 ⎯⎯→ [Al(NH2)]2+ + H+ + 3Cl–|
|In H2O:||: AlCl3 + H2O ⎯⎯→ [Al(OH)]2+ + H+ + 3Cl–|
|Amphoteric behaviour||In liq. NH3:
|Zn(NH2)2 + 2NH2– ⎯⎯→ [Zn(NH2)4]2–
Zn(OH2) + 2OH– ⎯⎯→ [Zn(OH)4]2–
Utility of the concept
Evidently this concept of solvent system can be used to explain the acid-base reactions occurring in aqueous solvents (protonic and non-protonic both).
- i) This theory does not consider a number of acid base reactions included in the protonic definition.
- ii) It limits acid base phenomena to solvent systems only. Thus it does not explain the acid – base reactions which may occur in the absence of solvent.
iii) It cannot explain the neutralisation reactions occurring without the presence of ions.
Thus this theory can simply be said to be an extension of the Arrhenius water – ion system.
Exercise 2 Write down the conjugate base of the following
(i) NH4+, (ii) HCOOH (iii) H3O+, (iv) H2NCONH3+
3.4 The Lewis Concept – The Electron – Donor – Acceptor System
This theory explains the acid-base phenomena not in terms of ionic reactions but in terms of electronic structure of the acid and base along with the formation of a coordinate covalent bond. According to Lewis (1923), an acid is any species (molecule, radical or ion) that can accept an electron-pair to form a coordinate covalent bond and a base is any species that can donate an electron-pair to the formation of a coordinate covalent bond. Thus, in the Lewis system, an acid is an electron pair-acceptor and a base is an electron pair-donor.
Thus according to Lewis theory, the process of neutralisation is simply the formation of a coordinate bond between an acid and a base. The neutralisation product, termed as coordinate complex or adduct, may be either non-ionisable or may undergo dissociation or condensation reaction depending on its stability.
Now consider the reaction between a proton (H+) and : NH3 molecules as shown below.
Evidently in the above reaction proton (H+) accepts one electron pair from :NH3 molecule and is, therefore, acid, whereas :NH3 molecule which donates an electron pair, is a base. The adduct is NH4+ ion.
Lewis bases and Bronsted-Lowry bases are the same substances, since any molecule or ion which accepts protons does so because of the presence of an unshared pair of electrons. In the above example NH3 molecules is a proton acceptor (i.e., Bronsted-base) and an electron pair donor (i.e., Lewis – base).
Bronsted and Lewis theories are thus identical as far as bases are concerned except that the wording used for definition of the bases is different in both the theories. Thus NH3, H2O, OH–, Cl–, CN– etc. are the bases on the Bronsted as well as on the Lewis systems. There are however, few compounds such as amides, ethers, nitriles, C2H4,C2H2,C6H6 etc. which have little or no tendency to accept protons but react readily with Lewis – acids.
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 from 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:
Lewis Acid + Lewis base ⎯⎯→ Adducts
- ii) Molecules containing a central atom with vacant d-orbitals. The central atom of the halides such as SiX4,GeX4,TiCl4,SnX4, PX3,PF3,SF4,SeF4,TeCl4 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 hydrolysed 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
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.
Lewis acid + Lewis base ⎯⎯→ Adduct or Addition compounds
Ammonation: Ag+ + 2(:NH3) ⎯⎯→ [NH3 ⎯⎯→ Ag ←⎯⎯ NH3]+
The Lewis acid strength or coordinating ability of the simple cation 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.
Evidently the acid strength of simple cations increases for the element on moving from left to right in a period and from bottom to top in a group of periodic table. Thus:
Fe2+ < Fe3+ (positive charge increases, +2 ⎯⎯→ +3)
K+ < Na+ (on moving from bottom to top in a group)
Li+ < Be2+ (on moving from left to right in a period)
⎯⎯⎯⎯⎯⎯⎯strength of Lewis acids increasing⎯⎯⎯⎯⎯⎯⎯→
- iv) Molecules having multiple bond 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 π-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 form dative bond.
SO2 also reacts in the same manner with OH– ion
- v) Elements with an electron sextent. Oxygen and sulphur atoms contain six electrons in their valence shell and can, therefore, be regarded as Lewis acids. The oxidation of to ion by oxygen and to ion by sulphur are the acid-base reactions.
|Lewis base||Lewis acid||Adduct|
Utility of Lewis concept:
- i) This concept also includes those reactions in which no protons are involved.
- ii) Lewis concept is more general than the Bronsted – Lowry concept (i.e. protonic concept) in that acid-base behaviour is not dependent on the presence of one particular element or on the presence or absence of a solvent.
iii) It explains the long accepted basic properties of metallic oxides and acidic properties of non-metallic oxides
- iv) This theory also includes many reactions such as gas phase, high temperature and non-solvent reaction as neutralisation process.
- v) The Lewis approach is, however, of great value in case where the protonic concept is inapplicable, for example, in reaction between acidic and basic oxides in the fused state.
- i) Since the strength of Lewis acids and bases is found to depend on the type of reaction, it is not possible to arrange them in any order of their relative strength. Thus, for example, experiments show that fluoride complex of Be2+ ions is more stable than that of Cu2+ ion, indicating that Be2+ ion is more acidic than Cu2+ On the other hand amine complex of Cu2+ is more stable than that of Be2+ ion indicating that Cu2+ is more acidic than Be2+ ion.
- ii) According to the phenomenological criteria, an acid-base reaction should be a rapid reaction. There are, however, many Lewis acid-base reactions which are slow.
Exercise 3: Among N,N dimethyl aniline and N,N,2,6 – tetramethyl aniline which one is a stronger base and why.
3.5 Hard and Soft Acids and Bases
The Lewis definition recognizes acids and bases in terms of their ability to accept or donate electron pairs. As such, the strength of any acid or base is determined by the very nature of the reaction involved in any particular electron transfer process. Accordingly, assignment of any single consistent criterion for acid base strength becomes very difficult in the Lewis definition. Attempts have been made to correlate various strength factors from enthalpy changes (ΔH) for acids – base combinations.
However, a qualitative correlation between the various Lewis acids and bases has been achieved. This is done by classifying the acids or bases into two classes hard and soft. Acid – base reactions are then treated by the general principle that hard acids prefer to combine with hard bases and soft acids perefer to combine with soft bases (and vice versa).
The criterion of hardness (or softness) is ascribed to the “hardness” of the electron cloud associated with any species. A firmly held electron–cloud with low polarisability makes a species “hard”, while an easily polarisable electron cloud characterises the species as “soft”. A third category with intermediate character appears in the border line. The principal distinguishing features of hard and soft acid and bases may then be summarised as follows:
|i) Small size||i) large size|
|ii) high positive oxidation state absence of any outer electron which are easily excited to higher states||ii) zero or low positive oxidation state several easily excitable valence electrons|
|H+, Li+, Na+, K+||Cu+, Ag+, Hg+|
|Be2+, Mg2+, Ca2+, Sr2+||Cd2+, Pt2+, Hg2+, Pt4|
|Al3+, Fe3+, Co3+, Cr3+||I+, I2, Br2, Br+|
|CO2,SO3||Cl, Br, I, N|
|HX (hydrogen –bonding molecules)||M° (metal atoms) and bulk metals|
|i) high electronegativity||i) low electronegativity|
|iii) Presence of filled orbitals; empty orbitals may exist at high energy level||ii) high polarisability (iii) partially filled orbitals; empty orbitals are low – lying.|
|H2O, OH–, F–, Cl–||H–, I–|
|, ,||SCN–, CN–,|
|, ,||R2S, RSH, RS–|
|NH3, N2H4, RNH2||R3P, R3As.|
Illustration 1: Arrange according to increasing Lewis acid character,
Solution: On the basis of the relative electronegativities of the halogens, one might expect the B atom in BF3 to be most electron deficient and hence most acidic. However, the compounds involve extensive π-interaction from a filled p orbital of the halogen and an empty p-orbital on B. The small F-atom forms the most stable p-p π-bond by efficient match of energy and size of the orbitals. Formation of an adduct with a base converts the geometry of bonds around boron from planar to pyramidal, thus rupturing the π-bonds. The Lewis acidity therefore increases with decrease in the extent of π-conjugation, as the halogen atom increases in size giving poorer overlap, i.e. from BF3 to BI3. Hence the order is
BI3< BF3 < BCl3 < BBr3 .
Illustration 2: Predict the relative acidic strength among the following:
HI, HIO4, ICl.
Solution: IO4– is more stable than I– and I+ is unstable. Hence HIO4 will be strongest acid.
HIO4 > HI> ICI
Illustration 3: Predict the relative acidic strength among the following :
H2O, H2S, H2Se, H2Te.
Solution: Assume that each has lost a proton : So, we get : HO–, HS– , HSe– , HTe–
It can be easily seen that the volume available for the negative charge is increasing from HO– to HTe–, Therefore :
- Volume available for the negative charge is increasing from left to right.
- Charge density is decreasing from left to right.
- Basicity is decreasing from left to right.
- Acidity of conjugate acids is increasing from left to right.
Illustration 4: Among maleic acid and fumaric acid which is a stronger acid and why.
Solution: Maleic acid is nothing but cis-butene dioic acid while fumaric acid is the trans isomer. The conjugate base of maleic acid i.e., the maleate ion is stabilized due to intramolecular hydrogen bonding which is not possible in the fumarate ion. Now more stable the conjugate base is the more more the acidity. Therefore maleic acid is stronger than fumaric acid.
Illustration 5: Arrange the following compounds in the increasing order of their basic strength.
CH3NH–Na+, C2H5NH2, (iso-C3H7)3N and CH3CONH2
Solution: CH3CONH2 < (iso-C3H7)3N, C2H5NH2, CH3ONH–Na+
Exercise 4: Arrange the following in increasing order of basicity
Phenol, o-nitrophenol, cresol
- Origin of Acidity and Basicity in Organic Compounds
- i) Among hydrocarbons, acidity increases with the % `s’ character . This is because, higher the `s’ character, closer are the electrons to the nucleus of C and farther the electrons go from H, easily can H be removed as H+.
Therefore CH3 – CH3 < CH2 = CH2 < CH ≡ CH
- Alcohols are more acidic than Alkanes because oxygen is more electro negative than C and consequently O – H bond easily gives H+ than C – H bond
H3C – H < CH3 – O – H
- Carboxylic acids are more acidic than alcohols because carboxylate anion is more resonance stabilised than carboxylic acid as compared to alkoxide anion with respect to alcohol.
Therefore CH3O – H < H -C – OH
- Phenols are more acidic than alcohols because phenoxide ion is more resonance stabilised as compared to phenol than alkoxide as compared to alcohol.
Therefore CH3 – O – H< C6H5 – O – H
- Phenols are less acidic than carboxylic acids. This is because in phenoxide ion some of the resonance structures are of higher energy and this makes phenoxide ion not very much stable as compared to phenol.
Therefore C6H5 O- H < H – C -OH
- In simple aliphatic acids, more the no. of CH3 added, less is the acidity. This is due to the +I (inductive) effect of CH3. This effect of adding CH3 decreases with increase in distance between CH3 added to the substituent, which releases H+.
CH3 -COOH > CH3 – CH2 – COOH
- Increase in `s’ character increases acidity of carboxylic acids.
CH3CH2COOH < CH2 = CHCOOH <HC ≡ C – COOH
- Increase in electron withdrawing substituents into simple aliphatic acids increases acidity.
CH3 – COOH < I – CH2 – COOH < Br – CH2 – COOH , < F – CH2 – COOH
- Nearer Cl is placed to COOH group, more is acidity.
- Benzoic acid is more acidic than carboxylic acids because benzoate anion shows/has more resonance.
CH3 COOH < C6H5 COOH
Exception to this is formic acid.
There is no destabilising factors for formate anion. Therefore, it is highly stabilised compared to formic acid because of resonance. Where as, Benzoic acid itself is resonance stabilised due to benzene ring. Therefore, benzoate ion is not very stable compared to benzoic acid. Thus C6H5 COOH < HCOOH
- More the no. of methyl or ethyl groups in NH3, more is its basicity. This is because of +I effect of methyl or ethyl group which helps nitrogen to donate its lone pair of electrons.
NH3 < CH3 – NH2 < (CH3)2 NH
NH3 < EtNH2 < (Et)2 NH
An exception to the above trend is the trisubstituted derivative. It is seen that if we introduce an alkyl group in a secondary amine, the basic strength of amines actually decreases. This is due to the fact that, the basic strength of an amine in water is determined not only by electron – availability on the nitrogen atom, but also by the extent to which the cation, formed by uptake of a proton, can undergo solvation, and so become stabilised. The more the number of hydrogen atoms attached to nitrogen in the cation, the greater the possibilities of powerful solvation via hydrogen bonding between these and water which increases stabilisation by solvation.
Thus, on going along the series NH3 → RNH2 → R2NH → R3N, the inductive effect will tend to increase the basicity, but progressively less stabilisation of the cation by hydration will occur, which will tend to decrease the basicity. The net effect of introducing successive alkyl groups thus becomes progressively smaller, and an actual change over takes place, on going from secondary to a tertiary amine. That is why, (CH3)3 N < (CH3)2 NH
If this is the explanation, no such changeover should be observed, if measurements of basicity are made in solvent in which hydrogen- bonding cannot take place; it has, indeed, been found that, in chlorobenzene the order of basicity of the butylamines is
BuNH2 < Bu2 NH<Bu3N
- Aniline is less basic than NH3. This is because in aniline the nitrogen atom is again bonded to an sp2 hybridised carbon atom but, more significantly, the unshared electron pair on nitrogen can interact with the delocalised orbital of the nucleus:
If aniline is protonated, any such interaction, with resultant stabilisation in the anilinium cation is prohibited, as the electron pair on N is no longer available :
The aniline molecule is thus stabilised with respect to the anilinium cation, and it is therefore energetically unprofitable for aniline to take up a proton ; it thus functions as a base with atmost reluctance.
- Introduction of alkyl, e.g. Me groups on the nitrogen atom of aniline results in small increase in pKa due to the +I effect of alkyl groups.
C6H5NH2 < C6H5 NHMe < C6H5NMe2
- Introduction of phenyl groups on N lowers basicity because the substituted amine becomes much more stable than the ion.
Ph2NH < PhNH2
Illustration 6: Why CHCl3 is more acidic than CHF3?
Solution: Cl3C–: is less basic than F3C–: because fluroine can disperse charge only by an inductive effect. while Cl (having empty 3d orbitals) disperses charge by inductive effect as well as by pπ – pπ bonding delocalisatioin. Fluorine is a second period element with no 2d orbital.
Illustration 7: Classify the following into acid, base and amphiprotic in terms of protonic concept.
- i) H2PO2–
- ii) H2PO3–
- iv) HPO3–2
- v) HPO43–
- vi) NH4–
Solution: Acidic (vi), (vii)
Basic (i) (iv)
Amphiprotic (ii), (iii), (v)
Illustration 8: SnO2 is an amphoteric oxide, explain
Solution: SnO2 + 2NaOH ⎯→ Na2SnO3 + H2O
SnO2 is acidic oxide since soluble is base
SnO2 + 4HCl ⎯→ SnCl4 + 2H2O
SnO2 is basic oxide since soluble in acid hence amphoteric base
Illustration 9: Give reason in one or two sentences for the following ammonium chloride is acidic in liquid ammonia solvent.
Solution: In solution of NH4Cl in liquid NH3, the following reaction takes place
NH4+ + NH3 NH3 + NH4+
Thus NH4Cl gives proton. Hence it is acidic
Illustration 10: Why PO43– ion is not amphiprotic?
Solution: An amphiprotic ion is one which can donate proton as well as accept proton. PO43– ion can accept proton but cannot donate any proton. Hence PO43– is not amphiprotic.
- Solution to Exercises
Exercise 1: The order in this case is the reverse of that for BX3. π-conjugation from the halogen p-orbital to the Si-d orbital is not as intense as in the case of BX3 and the order of acidity follows the increase in electron withdrawing power of the halogen from I to F. Hence the order is
SiI4 < SiBr4 < SiCl4 < SiF4
Exercise 2: To form a conjugate base means removal of a proton
So answer is
(i) NH3 (ii) HCOO–
(iii) H2O (iv) H2NCONH2
In N,N-2,6 – tetramethyl aniline the methyl groups on nitrogen and the ortho position are very close to each other resulting in a steric crowding. Now to avoid steric crowding the C—N bond rotates and becomes perpendicular to the benzene ring. In this process the lone pair on nitrogen becomes perpendicular to the p-orbitals of benzene ring thereby inhibiting resonance. But in N,N dimethyl aniline there is no steric hindrance, so the lone pair is in the same plane as the benzene ring and undergoes resonance. Therefore the lone pair on the tetramethyl derivative is more available and hence it is more basic.
Exercise 4: o-nitrophenol > phenol > cresol
- Assignments (Subjective Problems)
LEVEL – I
- Arrange in order of their acidity
H3O+, H2O, OH–
- Arrange according to increasing Lewis acid character,
- Arrange the decreasing order of acid strength
(I) CH2(OH)COOH; (II) Cl3COOH; (III) ClCH2COOH; (IV) CH3CCl2COOH
- Arrange in order of Lewis acid character of born trihalide will be Bi3, BBr3, BCl3, BF3.
- Between Na+ and Ag+ which is a stronger Lewis acid and why
- In dilute benzene solutions, equimolar additions of (C4H9)3N and HCl produce a substance with a dipole moment. In the same solvent, equimolar additions of (C4H9)3N and SO3 produce a substance having an almost identical dipole moment. What is the nature of the polar substances formed and what is the unifying feature of HCl and SO3?
- Arrange these acid in the increasing oder of their acidity.
HOCl, HO, HOBr, HOI
- Arrange according to increasing Lewis acid character,
SiF4, SiCl4, SiBr4, Sil4
9 Justify the statement that water behaves like an acid and also like a base on the basis of protonic concept.
- Predict the relative basic strength among the following :
LEVEL – II
1. Which of the following oxide is most acidic
Ag2O, V2O5, CO, N2O5
- Write down the order of conjugated base strength
OBr–, OCl–, OI–
- Arrange the following in correct order of acidity
- A certain reaction is catalysed by acids, and the catalytic activity for 0.1M solutions of the acids in water was found to decrease in the order HCl, HCOOH, CH3COOH. The same reaction takes place in anhydrous ammonia, but the three acids all have the same catalytic effect in 0.1M solutions. Explain.
- What is the order of increasing strength of α-chloro butanoic acid, β-chloro butnoic acid, γ-chloro butanoic acid n – butanoic acid.
- Write the cation, anion and the neutralisation reaction if thionyl chloride reacts with sodium sulphate.
- Arrange the following in the order of increasing acidity CH4, NH3, H2O & HF
- Among different hydroxy benzoic acids O-hydroxybenzoic acid is the strongest compared to m-and p-isomers considerably. Why?
- What should be the order of acidic strength in the series H3PO4, H3PO3 and H3PO2?
- Arrange the following in the increasing order of acidic strength.
LEVEL – III
- CCl4 does not act as a Lewis acid while SiCl4 and SnCl4 do so while SiCl4 and SnCl4 do so.
- BF3 act as a Lewis acid where as NF3 does not
- AsI2– complex is more stable than AsF2–.
- Arrange in their order of increasing acid strength.
HF, HCl, HBr, HI
- Increasing acid character
CO2, N2O5, SiO2, SO3
- Which way would the pH shift when CuSO4 is added to pure water. Write a net ionic equation to support your answer.
- Which one of the two i.e., CH3SH and CH2OH more strong acid and why?
- Compare the strength of formic acid and acetic acid
- Phenol is acidic but why ethanol is neutral
- Why F–– is most basic among the halogen anion
- Arrange the following in the increasing order of basicity. Justify anser
R3—N—, —R2NH, RNH2
- Among CH4, H2S, HI which one is most acidic? Justify your answer
- N-ethyl aniline is more basic than N methyl aniline why
- Arrange the following in the order of their basic strength. Justify your answer
- Arrange the following in the order of their decreasing acidic strength
CH4 , LiH, BeH2, H2O, HF, NH3
- Assignments (Objective Problems)
LEVEL – I
- Which one of following is the strongest acid
(A) H3PO4 (B) H3PO2
(C) H3PO3 (D) H3SO3
- Amongst the following the most basic compound is
- The conjugate acid of S2O82– is
(A) H2S2O8 (B) H2SO4
(C) HSO4– (D) HS2O8–
- What is the decreasing order of strength of bases
(A) CH3 – CH2− > NH2− > H – C ≡ C− > OH−
(B) H – C ≡ C− > CH3– CH2− > NH2− > OH−
(C) OH− > NH2− > H- C ≡ C− > CH3−CH2−
(D) NH2− > H – C ≡ C− > OH− > CH3 -CH2−
- Among the following weakest Lewis acid is
(A) H+ (B) OH–
(C) Cl– (D) HCO3–
- Which of the following is strongest Lewis acid
(A) CH3– (B) F–
(C) NH2– (D) OH–
- The strongest bronsted base in the following anions is
(A) ClO− (B) ClO2−
(C) ClO3− (D) ClO4−
- With reference to protonic acid which of the following statement is correct ?
- PH3 is more basic than NH3
- PH3 is less basic than NH3
- PH3 is equally basic than NH3
- PH3 is amphoteric while NH3 is basic
- Which of the following is a soft base
(A) R2S (B) NH3
(C) H2O (D) Cu=
- The conjugate base of CH3COOH is
(A) CH3COO– (B) CH3OH
(C) CH3COOH2+ (D) None of these
- The conjugate acid of S2O8–2
(A) H2S2O8 (B) H2SO4
(C) HSO4– (D) HS2O8–
- The conjugate base of H3O+ is
(A) H2O (B) OH–
(C) H+ (D) None of these
- Which of the following are amphiprotic in nature
(A) OH– (B)
(C) (D) HF
- Among the following compounds, the strongest acid is
(A) HC ≡ CH (B) C6H6
(C) C2H6 (D) CH3OH
- The following acids have been arranged in order of decreasing acid strength. Identify the correct order.
ClOH (I) BrOH (II) IOH (III)
(A) I > II > III (B) II > I > III
(C) III > II > I (D) I > III > II
LEVEL – II
- In which of the following interaction H2SO4 acts as a base?
(A) H2SO4 + HNO2 (B) H2SO4 + HNO3
(C) H2SO4 + HClO4 (D) None of these
- Which of the following is not Bronsted Lowry acid
(A) [Al(H2O)6]+3 (B) CH3COOH
(C) H2O (D) CH3NH2
- For AB3, Lewis base, the hybridisation state of A is sp3. It confirms the structure
(A) V-shape (B) Tetrahedral
(C) Square pyramidal (D) Trigonal pyramidal
- Weakest oxy-acid is
(A) HNO3 (B) H3PO4
(C) H3AsPO4 (D) H3SbO4
- The strongest acid is
(A) HClO3 (B) HClO2
(C) HClO4 (D) HClO
- The strength of an acid or base is determined by this process
(A) Reduction (B) Titration
(C) Hydration (D) Oxidation
- Bronsted acid is
- In which of the following cases the acid strength is the highest
(A) Ka = 10–6 (B) pKa = 5
(C) Kb = 10–11 (D) pKb = 10
- HF is a
(A) Strong oxidant (B) Strong reductant
(C) Weak acid (D) Strong acid
- Electron donor tendency is the highest far
(A) OH– (B) HS–
(C) HSe– (D) HTe–
- For the reaction of NH4 + S2– ⎯→ NH3 + HS–
(A) Acids (B) Bases
(C) Acid and pair (D) None of these
- Powerful Lewis acid character in which of the following cation
(A) Be+2 (B) Al3+
(C) Na+ (D) Li+
- Arrange the following acids in the increasing order of their acidity
(a) H3SO4+ (b) HSO4– (c) H2SO4
(A) b > c < a (B) b < c < a
(C) a > b > c (D) c > b > a
- Anhydrous HCl is a
(A) Acid (B) Base
(C) Salt (D) Covalent compound
- Weakest Bronsted base is
(A) H– (B) OH–
(C) Cl– (D) HCO3–
- Answers to Objective Assignments
LEVEL – I
- C 2. A
- D 4. A
- C 6. C
- A 8. B
- A 10. A
- D 12. A
- A 14. A
LEVEL – II
- C 2. D
- D 4. D
- C 6. B
- D 8. C
- C 10. A
- B 12. D
- B 14. D
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