Acids & Bases PDF

# Acids and Bases ## 19.1. Acid-Base Concepts - Introduction Acid and Base are very familiar terms, but it is difficult to define them exactly. However, they exhibit typical behaviour in chemical operations. Such behaviour may lead to their definitions which may be called operational definitions....

# Acids and Bases ## 19.1. Acid-Base Concepts - Introduction Acid and Base are very familiar terms, but it is difficult to define them exactly. However, they exhibit typical behaviour in chemical operations. Such behaviour may lead to their definitions which may be called operational definitions. The operational definitions of acids and bases were strictly experimental. An acid is a substance whose aqueous solution: - Turns blue litmus red. - Tastes sour. - Neutralises bases. - Reacts with active metals to liberate hydrogen. - Liberates CO2 from carbonates and bicarbonates. On the other hand, a base is defined as a substance whose aqueous solution: - Turns red litmus blue. - Tastes bitter. - Neutralises acids. - Gives a soapy touch. Some important theories of acids and bases are described below. ## 19.2. Arrhenius Concept of Acids and Bases Svante Arrhenius (1884) put forward his famous theory of electrolytic dissociation. According to this theory, an acid is a substance that dissociates in aqueous solution to give hydrogen ions. Thus, HCl, HNO3, H2SO4, CH3COOH and HCN are acids because they liberate H+ ions when dissolved in water. - HCl(g) + aq → H+(aq) + Cl¯(aq) - HNO3(l) + aq ←→ H+(aq) + NO3¯(aq) - H2SO4(l) + aq → 2H+(aq) + SO2 (aq) - HCN(g) + aq → H+(aq) + CN¯(aq) A base is a substance that dissociates in aqueous solution to give hydroxyl ions. Thus, NaOH, КОН, Ba(OH)2 and NH3 are bases because they liberate (OH¯) ions when dissolved in water. - NaOH(s) + aq → Na+(aq) + OH¯(aq) - KOH(s) + aq → K+(aq) + OH¯(aq) - Ba(OH)2(s) + aq → Ba2+(aq) + 2OH®(aq) - NH3(g) + aq NH‡(aq)+ OH¯(aq) ## 19.3. Advantages of Arrhenius Concept Arrhenius concept of acids and bases has the following applications or advantages. 1. It explains the acid-base neutralisation reactions easily. - H+ (acid) + OH¯ (base) → H2O (neutral). 2. In many reactions, H+ ions act as catalyst. For example, in the reaction CH3COOC2H5+H2O →CH3COOH + C2H5OH, H+ ions catalyse the hydrolysis of ethyl acetate to form acetic acid and ethyl alcohol. 3. It helps to express dissociation constant of acid (K) and base (Kg) in aqueous solution quantitatively. ## 19.4. Limitations of Arrhenius Theory 1. Limited Scope: Arrhenius concept is limited to aqueous medium only. It fails to explain the behaviour of acids and bases in non-aqueous solvents like ammonia, sulphur dioxide, alcohol etc. 2. Proton in aqueous solutions. According to Arrhenius concept, acids are those which furnish H+ ions in water. H+-ion is nothing but a proton, which is not capable of independent existence. It exists as hydronium ion i.e., H3O+. 3. Neutralisation without H+ and OH -ions: Neutralisation between HCl(g) and NH3(g) does not involve H⁺-ions and OH¯ ions e.g. - NH3(g) + HCl(g) → NH4Cl(s); CaO(s) + SO2(g) → CaSO3(s) ## 19.5. Bronsted-lowry Concept of Acids and Bases In order to overcome the limitations of Arrhenius concept, J.N. Bronsted and J.M. Lowry independently and simultaneously (1923) put forward a more general concept for acids and bases. This concept is called Bronsted-Lowry concept of acids and bases. According to this concept: - Acid is a substance (molecule or ion) which has a tendency to donate a proton to any other substance, and - A base is a substance (molecule or ion) which has a tendency to accept a proton from any other substance. In other words, (an acid is a proton donor and base is a proton acceptor). In fact, acid-base reaction involves transference of proton. That is why, this concept is also called proton transfer theory of acids and bases. In accordance with the broader definitions, an acid or base may be molecular or ionic, e.g. - **Molecular acids:** HF, HCl, HNO3, H2SO4, CH3COOH, H2S, H3PO4, H2O, HBr, HI, H2C2O4 (oxalic acid), H2CO3 (carbonic acid), NH2CONH2 (urea) etc. - **Ionic acids:** H3O+, NH†, HSO, HCO3, HPO¯, H₂SO₄, CH3 COOH, NH2CONH3, H2NO35 etc. - **Molecular bases:** RNH2, NH3, H2O, R2NH, H2SO4, CH3COOH, NH2CONH2(urea), HNO3 - **Ionic bases:** OH¯, §²¯, CO3¯², Cl¯, Br¯, NO3, F¯, ,I¯, HS¯, C202¯, SO, PO, HCO3, CH3COO¯, NO3 etc. It is important to note that no single substance is an acid or a base. A single substance cannot donate a proton unless and until some other substance which accepts the proton is also present. This proton transfer may or may not involve any medium, e.g. - NH3 + HCl → NH4 + + Cl does not involve any medium. Here, HCl has donated a proton to NH3. Hence, HCl is an acid with respect to NH3, which is a base. Consider some other acid-base reactions. - **NH3 + H2O** NH+ + OH¯ - **CH3COOH + NH3** → CH3COO¯ + NH‡ - **HCl + H2O** H3O+ + Cl¯ - **H2SO4 + H2O** 1 H3O+ + HSO A careful study of the above equations reveals that water is an acid with respect to NH3 whereas it is a base with respect to HCl or CH3COOH. Such substances like H2O, HCO3, HSO, etc. which can act as acids (proton donor) as well as bases (proton acceptor) are called amphoteric or amphiprotic substances. **Conjugate Acid-Base pairs:** When an acid loses a proton, the residual part of it will have a tendency to accept a proton. Thus, it will behave as a base. Such pairs of substances which differ from one another by proton are known as conjugate acid-base pairs. Consider a general example of an acid (HA) - **HA → H+ + A¯** Consider the reaction between HCl and H2O. - **HCl + H2O** H3O+of HCl donates a proton to water in the forward reaction. - HCl is an Acid and, H2O is a Base. In the backward reaction, H3O+ ion donates its proton to Clion. - H3O+ is an acid and, Cl is a base. It is clear that HCl loses a proton and forms Cl-ion (which is a base). HCl and Cl¯ ion are called conjugate acid-base pair. Similarly, H2O (base) accepts a proton and forms H3O+ (acid). Here, H2O and H3O+ are also conjugate base-acid pair. - **H2O + H+** H3O+ In the overall reaction between, HCl and H2O, can, hence, be written as: - **HCl + H2O = H3O+ + CI** Such reactions are called neutralisation reactions. ## 19.6. Relative Strength of Acids and Bases According to proton transfer theory (i) the strength of an acid depends upon its tendency to donate a proton and (ii) strength of a base depends upon its tendency to accept a proton. For example, consider the reaction between acetic acid and water, - **CH3COOH + H2O → CH3COO¯ + H3O+** Acetic acid has a small tendency to donate a proton. Therefore, it is a weak acid. The above equilibrium lies mostly towards left hand side. It follows, therefore, that CH3COO¯ ion must have a strong tendency to accept a proton. Hence, acetate ion (conjugate) base of CH3COOH is a strong base. - **Weak acid H+ + Strong conjugate base.** Similarly, HCl is a strong acid in water. - **HCl + H2O H3O+ + Cl¯** The above equilibrium lies mostly towards right hand side. It follows therefore, that Cl¯ ion must have a little tendency to accept a proton. Hence, Cl¯ ion is a weak base. - **Strong Acid H⁺ + weak conjugate base** Hence, we conclude that every strong acid has a weak conjugate base and vice versa. ## 19.7. Amphiprotic Substances Substances (molecules or ions) which can act both as Bronsted acids (proton donor) and Bronsted bases (proton acceptor) are called amphiprotic substances. For example, urea (NH2CONH2), HS, HSO₄, HNO3, CH3COOH, H2SO4, H2O etc. The acidic and basic nature of above amphoteric substances are given below : | | Base | Acid | Acid2 | Base2 | |---|---|---|---|---| | (a) | NH3 | NH2CONH2 | NH | NH2CONH | | | | HCIO4 | NH2CONH | CIO | | (b) |NH3 | HS | NH | $2- | | | | H2O | H2S | OH | | (c) | H2O | HSO | H3O+ | so²- | | | | HSO | H2SO4 | H2O | | (d)| NH3 | HNO3 | NH | NO3 | | | | HNO3 | HF | F | | | | | H2NO3 | | | (e) | NH3 | CH3COOH | NH | CH3COO¯ | | | | | CH3COOH | | | (f) | NH3 | H2SO4 | 2NH | SO²- | | | | H2SO4 | HF | F | | | | | H3SO | | | (g) |NH3 | H2O | NH | OH | | | | | H3O+ | NO3 | ## 19.8. Applications of Amphoteric Nature of Water The amphoteric nature of water helps to explain the acidic or basic nature of some salts on the basis of hydrolysis. For example: - **(i) Aqueous solution of copper (II) sulphate is acidic in nature.** When copper (II) sulphate is dissolved in water, it ionises as follows : - CuSO4 Cu2+ + So The SO ions formed above have a stronger tendency than H2O to accept a proton. So, SO ions accept H+ ions from water to form highly ionised sulphuric acid. As a result, excess of H+ ions are available in the solution. So, CuSO4 solution is acidic in aqueous solution and turns blue litmus red. - 2H2O(acid) + SO2 (base) → 2H+ + SO4 + 2OH - Cu2+ + 2OH → Cu(OH)2↓ - **(ii) Aqueous solution of Na2CO3 is basic in nature.** When Na2CO3 (sodium carbonate) is dissolved in water, it ionises as follows : - Na2CO3 2Na+ + CO3- The CO3 ions formed above have a stronger tendency than H₂O to accept a proton. So, CO3¯ ions accept H+ ions from water to form weakly ionised carbonic acid (H2CO3). As a result, excess of OH ions are available due to highly ionised NaOH formed as shown below. Hence, Na2CO3 solution is basic in nature and turns red litmus blue. - 2Na+ + CO3 + 2H2O2Na+ + 2OH + H2CO3 - **(iii) Aqueous solutions of ferric (III) and altiminium (III) salts are acidic in nature.** Fe3+ and Al3+ salts exist as hexahydrated cations [Fe(H2O)6]³+ and [Al(H2O)6]³+ in aqueous solution. These cations have greater tendency to donate a proton than H2O. As a result, H3O+ (acid ions) ions are formed. So, their aqueous solution is acidit and turn blue litmus red. - **[M(H2O)6]³+ (M=Fe, Al) (acid)+H2O (base) → H3O+ (acid)+[M(H2O)5(OH)]²+(base)** ## 19.9. Monoprotonic and Polyprotonic Bronsted Acids and Bases **Monoprotonic Bronsted acids** are those substances which lose only one proton. For example, - HCl ; (HCl → H+ + Cl¯) - CH3COOH ; (CH3COOH → CH3COO¯ + H+) ; HCN → H+ + CN¯ **Polyprotonic Bronsted acids** are those substances which lose two or more protons. For example, oxalic acid, carbonic acid (H2CO3) etc. - H2CO3; (H2CO3 → 2H+ + CO¯), H₂SO₄ → 2H+ + SO - H2S 2H+ + S²- **Monoprotonic Bronsted bases** are those substances which can accept one proton. For example, water. - (H2O + H+ → H3O+), HSO3; (HSO3 + H+ → H2SO3) - HS¯; (HS¯ + H+ → H2S), - NH2CONH¯ (NH2CONH¯ + H+ → NH2CONH2, urea), - NH3 + H+ → NH. **Polyprotonic Bronsted Bases** are those substances which can accept two or more protons. For example, - SO¯; (SO¯+2H+ → H2SO4) - PO3; (PO3 + 3H+ → H3PO4) - C2O4 oxalate ion; (C₂O¯ + 2H+ → H2C2O4) ## 19.10. Comparison with Arrhenius Concept Arrhenius 'concept as well as Bronsted-Lowry concept consider an acid as a source of protons. For example, HCl is an acid according to both the concepts. In general, we can say that all Arrhenius acids are also Bronsted acids, but reverse may not be true. For example, NH ion is an acid according to Bronsted theory but not so according to Arrhenius. All Bronsted bases may not be Arrhenius bases Reason. According to Arrhenius, base must furnish OH¯-ions in aqueous solution. According to Bronsted theory, a base must accept a proton. Example. According to Bronsted theory Co ion is a base as it accepts a proton to form HCO3 ion. - **CO3¯ + H+ → HCO3** According to Arrhenius concept, CO3 is not a base as it does not furnish OH¯-ions in aqueous solution. Hence, all Arrhenius acids are also Bronsted acids but all Bronsted bases are not Arrhenius bases. ## 19.11. Limitations of Bronsted Lowry Theory Bronsted and Lowry's concept has a wider view. It covers molecular as well as ionic species. This concept does not require any medium like H₂O, as needed in Arrhenius theory. However, it has got the following limitations : 1. It fails to explain the reaction between acidic oxides such as CO2, SO2, NO2 and basic oxides such as CaO, BaO, MgO etc. 2. It fails to explain the acidic character of AlCl3, FeCl3, BF3 etc. ## 19.12. Lux-flood Concept It is non-protonic concept of acid-base reactions. Lux (1939) observed that acid base reactions are also possible in oxide systems without the help of protons. Flood (1947) extended the Lux-concept and applied it to the non-protonic systems which were not convered by Bronsted and Lowry. According to this concept : - An acid is that substance which accepts an oxide ion while a base is that substance which donates (or gives up) an oxide ion. Following are some acid-base reactions in terms of Lux-Flood Concept. General reaction. - **Base O²¯ + Acid → O2 + 2Na+** Typical reactions - **Na2O → O²¯ + 2Na+** - **BaO → O²¯ + Ba²+; SO2 → O2 + SO3** Above reactions indicate that the base is an oxide donor while acid is an oxide acceptor. Based on the above definitions, a few direct reactions between acidic. and basic anhydrides in the absence of hydrogen ions or water are given below : | Base | Acid | |---|---| | 6Na2O | P4010 | | CaO | SiO2 | | BaO | CO2 | | Na2O | ZnO | - **Such direct reactions are in conformity with the following reaction sequence** - BaO + H2O → Ba(OH)2(aq) (Base); CO2 + H2O → H2CO3 (aq) (acid) **Scope-Applications.** - **(a) This view is particularly applicable to reactions which take place at high temperature, i.e. reactions during the manufacture of glass, ceramics, and in metallurgical operations.** Thus, the ores of titanium (Ti), tantalum (Ta) and niobium (Nb) can be brought into solution at 1073 K in sodium pyrosulphate (Na2S2O7) or potassium pyrosulphate (K2S2O7). - K2S2O7 + TiO2→ K2SO4 + (TiO) SO4 or Ti(SO4)2 **Some other examples involving above type of reactions are :** - Δ - PbO (base) + SO3 (acid) → PbSO4; CaO (Base) + SiO2 (acid) → CaSiO3 - **(b) This concept also classifies substances as amphoteric. Amphoteric substances are those which show both a tendency to take up or give up oxide ions depending on the circumstances.** - For example : ZnO + Na2O → 2Na+ + ZnO½¯; ZnO + S2O3¯ → Zn2+ + 2SO¦¯ - **(c) The Lux-Flood oxide-transfer picture of acid-base reactions can be extended to any negative ion i.e., sulphides, halides or even carbanions. Some examples are given for clarity.** | Negative ion donor (base) | Negative ion acceptor (acid) | |---|---| | 3NaF | AlF3 | | C2H5Na | (C2H5)2Zn | | Na2S | CS2 | - **(d) According to Flood and Forenson, a parallelism, between the thermodynamic stabilities of carbonates and sulphates to-wards the evolution of CO2 and SO2 gases respectively at high temperature also exists.** The order of increasing ease of decomposition of carbonates and sulphates is - Ba2+>Li+>Ca² > Mg2+ > Mn2+ > Cd2+ >Pd²+ > Co²+ > Ag+ > Fe2+ >Ni2+>Cu²+>Fe3+>Be2+. This order is nearly the same as the order of basic strength of the hydroxides of these metals. **Limitations:** This concept is applicable to only one part of the general theories of Usanovich and Lewis. ## 19.13. Usanovich's Concept (or Positive-Negative System) M. Usanovich, a Russian chemist proposed a new concept of acids and bases (in 1939), called Usanovich's concept of acids and bases. According to this concept: - **(A)** An acid is any chemical species which : - (i) reacts with (or neutralises) a base to form a salt. - (ii) accepts electrons or anions and (iii) furnishes cations. - **(B)** A base is any chemical species which : - (i) reacts with (or neutralises) an acid to form a salt - (ii) accepts cations and (iii) furnishes electrons or anions. Following examples are given for clarity 1. **Action of SO2 (acid) on Na2O (base) to form Na2SO3 (salt).** - SO2 (acid) + Na2O (base) → Na2SO3 (salt) **Explanation.** - **(a)** - Na2O→ 2Na+ + 02- - **Since Na2O furnishes O²¯ anion, it is a base.** - **(b)** - SO2 + O²¯→ SO3 - **Since SO2 has accepted the anion, O2, it is an acid.** Adding equations (ii) and (iii), we get equation (i). - **Na2O + SO2 + O2- → 2Na+ + O²¯ + SO3¯ → Na2SO3** 2. **Action of SO3 (acid) on Na2O (base) to form Na2SO4 (salt)** - SO3 (acid) + Na2O (base) → Na2SO4 (salt) **Explanation.** Na2O is a base because it furnishes the anion, O²¯ (Na2O → 2Na+ + O2¯). SO3 is an acid because it accepts the anion, O2 (SO3 + 0²¯ → SO²¯). 3. **Action of Fe(CN)2 (acid) on KCN (base) to form K4 [Fe(CN)6] (salt)** - Fe(CN)2 (acid) + 4 KCN (base) → K4[Fe(CN)6] (salt) **Explanation.** KCN is a base because it furnishes the anion, CN¯(KCN→ K+ + CN¯). Fe (CN)2 is an acid because it accepts the anion, CN¯ {Fe(CN)2 + 4CN® → [Fe(CN)6]^¯} 4. **Cl (acid) reacts with Na (base) to form NaCl (salt).** - Cl2 (acid) + 2 Na (base) → 2NaCl (salt) **Explanation.** - **2Na → 2Na+ +2e** - **Since Na furnishes an electron, it is a base.** - **2Cl + 2e- 2C1** - **Since Cl has accepted an electron, it is an acid.** Adding equations (ii) and (iii), we get equation (i). - **2Na + 2Cl + 2e- → 2Na+ + 2e+ 2Cl2NaCl** It is a redox reaction also because in this reaction, Na is oxidised to Na⁺ and Cl is reduced to Ct. 5. **Sb2S5 (acid) reacts with (NH4)2 S (base) to form (NH4)3 SbS4 (salt)** - Sb2S5 (acid) + 3 (NH4)2S (base) → 2 (NH4)3 SbS4 **Explanation.** - **(a)** - (NH4)2S → 2NH+ + S²- - **Since (NH4)2S furnishes the anion, S2, , it is a base.** - **(b)** - Sb2S5 + 3S22SbS3- - **Since Sb2S5 accepts the anion, S2, it acts as an acid.** Adding equations (ii) and (iii), we get equation (i) - **(NH4)2S → 2NH + S²-1 × 3** - **Sb2S5 + 3S22SbS** - **3(NH4)2S + Sb2S5→6NH4+ 2SbS→ 2 (NH4)3 SbS4** 6. **Reaction between CO2 (acid, O = C = O) containing unsaturated carbon, and OH (base) to form bicarbonate (HCO3 or H – O – C – O¯) was also considered by Usanovich.** - O = C = O (acid) + OH¯ (base) → HCO3 Usanovich emphasises on the ability of the central atom in the compound (acid) to have co-ordination unsaturation (i.e., ability of the central atom to increase its covalency by accepting electrons into its vacant orbitals). Thus, the acidic function of a compound is determined by the presence of coordinately unsaturated positive particles. According to him, the basic function of a compound is determined by the presence of unsaturated negative particles. For example, the central atom 'S' in SO2 is coordinately unsaturated due to the presence of vacant 3d-orbitals. So, it accepts the electrons of O² anion to form SO3 and hence it acts as an acid. Here 0²- anion acts as a base. The above explanation is similar to that of Lewis acid-base concept. **Advantages (or merits).** - **1.** It does not lay emphasis on the donation or acceptance of electrons in the form of shared pairs. - **2.** It covers all the acids and bases defined by Lewis. - **3.** This concept is like the synthesis of all the previous theories of acids and bases and has, thus, become the most general definition. - **4.** Even the redox reactions have been termed as acid-base reactions. For example, consider the following reaction: - **2Na + Cl2→ 2NaCl** . This reaction can be explained as follows : - **Na → Na⁺ +e¯] × 2[Na is oxidised to Na⁺. So it is oxidation reaction]** - **Cl2 + 2e- → 2Cl¯ ] [Cl2 is reduced to Cl¯. So, it is a reduction reaction]** - **2Na + Cl2→ 2Na+ + 2Cl2 NaCl (redox reaction)** **Limitations**. - **1.** It is so general that it cannot be of wide applicability. - **2.** It is so general that atmost all chemical reactions can be considered as acid-base reactions. ## 19.14. Lewis Concept or Electronic Concept ## 19.15. Effect of Substituents on the Strength of Lewis Acids and Bases The effect of substituent on the strength of Lewis acid and Lewis base depends upon the nature of substituent attached to the acceptor atom of Lewis acid and donor atom of Lewis base. In order to compare the relative strength of acids and bases the concept of inductive, resonance and steric effects is useful. - **1. Inductive effect:** When an electron withdrawing group (F, Cl, Br, I, NO2 etc.) is attached to the donor atom of a Lewis base, it withdraws some electron density from the donor atom and thereby decreases the electron density on it. Consequently, the tendency of the atom to donate the electron pair decreases and hence strength of the Lewis bases decreases. When an electron withdrawing group is attached to the acceptor atom of a Lewis acid, it withdraws some electron density from the acceptor atom and thereby decreases the electron density on it. Consequently the tendency of the atom to accept electron pair from bases increases and hence strength of Lewis acids lecreases. In general, "an electron withdrawing group increases the acid strength and decreases the basic strength of Lewis acids and Lewis bases respectively." When an electron donating group is attached to the Lewis acid or Lewis base, an opposite effect to that explained above is observed i.e., "an electron releasing group decreases the acid strength and increases the basic strength of Lewis acids and Lewis bases respectively." When an electron releasing group is attached to the donor atom of Lewis bases, it increases the electron density on the donor atom. As a result, its basic character (i.e., tendency to donate electrons) increases. When an electron releasing group is attached to the acceptor atom of Lewis acid, it increases the electron density on acceptor atom. As a result, its acidic character (i.e., tendency to accept electron) decreases. As a result, its acid strength decreases. Consider BH3 molecule. It is a Lewis acid because it is deficient in electrons i.e. it has only six electrons in its valence shell which are two electrons short from octet. If H-atoms are replaced by electron releasing groups (-CH3), the acidic strength decreases. If the H-atoms are replaced by electron withdrawing groups (F), the acidic strength increases. Thus the increasing order of acidic strength of BH3, (CH3)3B and BF3 is (CH3)3B < BH3 < BF3. Comparison of basic strength of NH3, NF3 and (CH3)3N. All these compounds are basic because of lone pair of electrons on their central atom, nitrogen. The methyl group is an electron releasing group. So, it will increase the electron density on N-atom of (CH3)3N. As a result, (CH3)3N becomes more basic than NH3. The F-atom is an electron withdrawing group. So, it will decrease the electron density on N-atom of NF3. As a result, NF3 becomes less basic than NH3 (fig. 19.1). Hence the decreasing order of basic strength of (CH3)3N, NH3 and NF3 is (CH3)3N > NH3 > NF3 **2. Resonance effect:** The relative acid strength of halides of boron (BF3, BCl3 and BBr3) can be explained with the help of resonance effect. Boron halides are converted to their adducts with pyridine. Then, the heats of formation and dipolemoments of these adducts are measured. From the data, it is observed that the relative electron pair acceptor strength of these halides is in the following order : - BBr3 > BCl3 > BF3. On the basis of steric concept and electronegativity concept of F, Cl and Br, the relative electron pair acceptor strength of these halides should be in the order. - BF3 > BC13 > BBr3 **Explanation of anomalous behaviour. **The above trend in the acidic strength of boron halides can be explained in terms of boron-halogen π-back bonding. F-atom has filled 2p-orbital while B-atom has vacant p-orbital of same energy. The filled 2p-orbital of F-atom overlaps effectively with vacant 2p-orbital of B-atom. As a result, pr – pr dative or back bond is formed. This B-F bond has some double bond character. The various resonating structures of BF3 are given in figure 19.2. Due to back bonding (Fig. 19.3) electron deficiency of B-atom is compensated to some extent. As a result, lewis acidity of BF3 gets decreased. The tendency of back bonding decreases in the order BF3 > BCl3 > BBr3. It is because the overlap of vacant 2p-orbital (lower energy) with filled 3p and 4p-orbitals (of higher energy) does not take place effectively. Hence the relative Lewis acid strength of BF3, BCl3 and BBr3 is as follows : - BBr3 > BC13 > BF3 **Relative Lewis acid strength of halides of group 13 elements.** As we go down the group, the size of atoms goes on increasing. As a result, their electron pair acceptor ability (or Lewis acidity) goes on decreasing in the order B > Al > Ga > In. On the basis of resonance concept we can similarly explain that (CH2O)3B is a weaker Lewis acid than trimethyl boron, (CH3)3B. **3. Steric Effect:** The relative strength of Lewis acids and bases is also affected by the bulky groups present in them. When a Lewis acid (say boron or boron halide) is treated with a Lewis base (say amine), the bulky groups present in Lewis acid and/or base affect the stability of acid-base adduct and hence their strength. Based on above, let us compare the basic strength of pyridine, 2-methyl pyridine (2-or a-picoline) and 4-methyl pyridine (4- or y-picoline). Methyl group is electron releasing group. So, because of its positive inductive effect, 2-methyl pyridine and 4-methyl pyridine are expected to be stronger bases than pyridine. The basicity of these amines with respect to trimethyl boron, B(CH3)3 is of the following order. - pyridine ≈ 4-methyl pyridine > 2-methyl pyridine The least basic character of 2-methyl pyridine is because of the steric hindrance between CH3 group in B(CH3)3 and CH3 group in 2-methyl pyridine. The effect is called front or F-strain (fig. 19.4). The CH3 group in 2-methyl pyridine is very near to the bulkier CH3 groups of B (CH3)3 and their adduct is unstable (∆H = 42 kJ mol¯¹). The adducts with pyridine and 4-methyl pyridine are stable as clear from their ΔΗ values. Similarly, we can explain that quinuclidine is stronger base to-wards trimethyl borane than triethylamine. In triethyl amine, there are three bulky ethyl groups which surround the nitrogen atom and cause steric hindrance (fig. 19.5). The above fact of basic strength is supported by the bond energy of quinuclidine (CH3)3B adduct (= 84 kJ mol) which is more than the bond energy of quinuclidine – (C2H5)3 B adduct (= 42 kJ mol¯). It is interesting to note that 2-tert. butyl pyridine (Lewis base) does not form complex with trimethyl borane although a lone pair of electrons is present on the N-atom of the Lewis base. It is due to steric hindrance. ## 19.16. Comparison with Bronsted Theory Every Bronsted base is also a Lewis base. Bronsted base has got the ability to accept a proton from an acid. Now, any substance that can accept a proton may donate a pair of electrons e.g. NH3 can accept a proton. - : NH3 + H+ → NH NH3 has got a lone pair to share with H⁺. Thus, NH3 is Lewis base. All Bronsted acids are not the Lewis acids. For example, CH3COOH is a Bronsted acid because it can donate a proton to : NH3 to form ammonium ion. - CH3COOH + NH3 → CH3COO¯ + NH However, it is not a Lewis acid because it is not deficient in electrons. The proton that can be donated (to some base) by acetic acid is the Lewis acid. ## 19.17. Limitations of Lewis Concept - This concept has following limitations. - **1.** It is too general and covers the formation of all co-ordination compounds. - **2.** Common acids and bases do not form co-ordinate bonds, whereas co-ordinate bond formation is an essential feature of Lewis concept. - **3.** It does not help to explain the relative strengths of acids and bases. - **4.** Process of neutralisation between common acids and bases is instantaneous whereas co-ordinate bond formation is a slow process. - **5.** It