Ionic Equilibrium PDF

60 340 Ionic Equilibrium E3 Chapter 9 Ionic Equilibrium Electrical conductors For example : H2O, CH 3COOH , NH 4 OH, HCN, Liq. SO 2 , HCOOH etc. means all weak acids and bases. U compound then it reacts with water to form ions. Degree of ionization for this types of electrolytes is   1. ID In ch...

60 340 Ionic Equilibrium E3 Chapter 9 Ionic Equilibrium Electrical conductors For example : H2O, CH 3COOH , NH 4 OH, HCN, Liq. SO 2 , HCOOH etc. means all weak acids and bases. U compound then it reacts with water to form ions. Degree of ionization for this types of electrolytes is   1. ID In chemical equilibrium we studied reaction involving molecules only but in ionic equilibrium we will study reversible reactions involving formation of ions in water. When solute is polar covalent D YG Substances, which allow electric current to pass through them, are known as conductors or electrical conductors. Conductors can be divided into two types, (1) Conductors which conduct electricity without undergoing any chemical change are known as metallic or electronic conductors. Arrhenius theory of electrolytic dissociation (1) Postulates of Arrhenius theory (i) In aqueous solution, the molecules of an electrolyte undergo spontaneous dissociation to form positive and negative ions. (ii) Degree of ionization ( )  Number of dissociate d molecules Total number of molecules of electrolyt e before dissociati on (iii) At moderate concentrations, there exists an equilibrium between the ions and undissociated electrolytes. molecules, such as, NaOH ⇌ Na   OH  ; KCl ⇌ K  U (2) Conductors which undergo decomposition (a chemical change) when an electric current is passed through them are known as electrolytic conductors or ST Electrolytes are further divided into two types on the basis of their strengths, (i) Substances which almost completely ionize into ions in their aqueous solution are called strong electrolytes. Degree of ionization for this type of electrolyte is one i.e.,   1.  Cl  This equilibrium state is called ionic equilibrium. (iv) Each ion behaves osmotically as a molecule. (2) Factors affecting degree of ionisation (i) At normal dilution, value of  is nearly 1 for For example : HCl, H 2 SO 4 , NaCl , HNO 3 , KOH , NaOH , strong electrolytes, while it is very less than 1 for weak electrolytes. HNO3 , AgNO 3 , CuSO 4 etc. means all strong acids, bases (ii) Higher the dielectric constant of a solvent more is its ionising power. Water is the most powerful and all types of salts. (ii) Substances which ionize to a small extent in their aqueous solution are known as weak electrolytes. ionising solvent as its dielectric constant is highest. Ionic Equilibrium (iii)   1 Con. of solution  1 wt. of solution (iv) Degree of ionisation of an electrolyte in solution increases with rise in temperature. (v) Presence of common ion : The degree of ionisation of an electrolyte decreases in the presence of Furher, if V is the volume of the solution in litres containing 1 mole of the electrolyte, C  1 / V. Hence we have   Ka V in water, it dissociates partially into H  or H 3 O  and following equilibrium is CH 3COOH  H 2O ⇌ CH 3 COO  H 3 O  “For a weak electrolyte, the degree of ionisation is inversely proportional to the square root of molar concentration or directly proportional to the square root of volume containing one mole of the solute.” Dissociation constants of acids and Bases (1) Dissociation constant for weak acid : Consider an acid HA which, when dissolved in water ionizes as, Applying law of chemical equilibrium, [CH 3 COO  ]  [H 3 O  ] [CH 3 COOH ]  [H 2O] HA ⇌ H   A  U In dilute solution, [H 2 O] is constant. The product K The above equations lead to the following result ID  of …..(v) This is called Ostwald’s dilution law. obtained, K 60 When acetic acid (a weak electrolyte) is dissolved   Kb / C  Kb V E3 The strength of an acid or a bas is experimentally measured by determining its dissociation or ionisation constant. the …..(iv) Similarly, for a weak base like NH 4 OH , we have Ostwald's dilution law ions and …..(iii) The degree of dissociation,  can therefore be calcualted at a given concentration, C if K a is known. a strong electrolyte having a common ion. CH 3 COO  Ka C K a  C  2 or    Dilution of solution  Amount of solvent 341 and constant [H 2 O] is denoted as K a , the Ka  D YG ionization constant or dissociation constant of the acid is, [CH 3 COO  ]  [H 3 O  ] [CH 3 COOH ] …..(i) The fraction of total number of molecules of an electrolyte which ionise into ions is known as degree of dissociation/ionisation . If ' C' represents the initial concentration of the acid in moles L1 and  the degree of dissociation, U then equilibrium concentration of the ions (CH 3 COO  and H 3 O  ) is equal to C  and that of the undissociated acetic acid  C(1   ) i.e., we have ST C Conc. at eqb. C(1   ) 0 C Substituting the values of the concentrations in equation (i), we get Ka  C .C  C 2 2 C 2   C(1   ) C(1   ) 1   0 C equilibrium Where, K a is the dissociation constant of the acid, HA. It has constant value at definite temperature and does not change with the change of concentration. Dissociation Constant for polybasic acid : Polybasic acids ionise stepwise as, for example, orthophosphoric acid ionises in three steps and each step has its own ionisation constant. H 3 PO 4 ⇌ H   H 2 PO4 (I step) H 2 PO 4 ⇌ H   HPO 42 (II step) HPO 42 (III step)  ⇌ H  PO 43 Let K1 , K 2 and K 3 be the ionization constants of K1  [H  ][PO 43 ] [H  ][HPO 42 ] [H  ][H 2 PO 4 ] K  ; K2  ; 3 [H 3 PO 4 ] [HPO 4 2 ] [H 2 PO 4 ] In general, K1  K 2  K 3 The overall dissociation constant (K ) is given by …..(ii) In case of weak electrolytes, the value of  is very small and can be neglected in comparison to 1 i.e., 1   1. Hence, we get [H  ][ A  ] [HA ] first, second and third steps respectively. Thus, CH 3COOH  H 2O ⇌ CH 3 COO   H 3 O  Initial conc Applying the law of mass action, K a  the relation, K  K1  K2  K3 (2) Dissociation constant for weak base : The equilibrium of NH 4 OH (a weak base) can be represented as, NH 4 OH ⇌ NH 4  OH  342 Ionic Equilibrium K b is constant at a definite temperature and does not change with the change of concentration. Common ion effect The degree of dissociation of an electrolyte (weak) is suppressed by the addition of another electrolyte (strong) containing a common ion, this is termed as common ion effect. Acetic acid is a weak electrolyte and its ionisation is suppressed in presence of a strong acid ( H  ion as common ion) or a strong salt like sodium acetate (acetate ion is a common ion). Similarly, the addition of NH 4 Cl or NaOH to NH 4 OH solution will Applying the law of mass action, [ A  ][B  ] K [ AB] Since the solution is saturated, the concentration of unionised molecules of the electrolyte is constant at a particular temperature, i.e., [ AB]  K   constant. Hence, [ A  ][B  ]  K[ AB ]  KK   K sp (constant) K sp is termed as the solubility product. It is defined as the product of the concentration of ions in a saturated solution of an electrolyte at a given temperature. Consider, in general, the electrolyte of the type 60 [ NH 4 ][OH  ] [ NH 4 OH ] Applying the law of mass action, K b  A x By which dissociates as, A x By ⇌ xA y   yB x  [ A y  ]x [B x  ]y K [ A x By ] Applying law of mass action, either NH 4 or OH . When the solution is saturated, NH 4 OH ⇌ NH 4 CH 3 COONa CH 3 COO   Na  NH 4 Cl NH 4  H   OH  Cl  Common ion Common ion  adjust concentration of S 2  ions in second group and OH  ion concentration in third group. D YG Isohydric solution If the concentration of the common ions in the solution of two electrolytes , for example in HCl and HNO 3 or H OH  ion ion concentration in Ca(OH )2 and Ba(OH )2 is same, then on mixing them there is no change in the degree of dissociation of either of the electrolytes. Such solutions are called isohydric solutions. U Consider two isohydric solutions of acids HA1 and HA2. Let V1 and V2 be their dilutions and 1 and  2 be their degree of dissociation at the respective ST dilution. Then, 1 V1  (constant) or [ A y  ]x [B x  ]y  K[ Ax By ]  KK   K sp (constant) Thus, solubility product is defined as the product of concentrations of the ions raised to a power equal to the number of times the ions occur in the equation representing the dissociation of the electrolyte at a given temperature when the solution is saturated. (1) Difference between solubility product and ionic product : Both ionic product and solubility product represent the product of the concentrations of the ions in the solution. The term ionic product has a broad meaning since, it is applicable to all types of solutions, either unsaturated or saturated and varies accordingly. On the other hand, the term solubility product is applied only to a saturated solution in which there exists a dynamic equilibrium between the undissolved salt and the ions present in solution. Thus the solubility product is in fact the ionic product for a saturated solution at a constant temperature. (2) Different expression for solubility products (i) Electrolyte of type AB (1 : 1 type salt) e.g., AgCl , BaSO 4 U As a result of common ion effect, the concentration of the ion of weak electrolyte which is not common in two electrolytes, is decreased. The use of this phenomenon is made in qualitative analysis to concentration 2 AgCl ⇌ Ag   Cl  x x V2 Ksp  [ Ag  ][Cl  ] ; Ksp  x 2 ; x  K sp Above equation is useful for calculating the relative dilution of two acids at which they would be isohydric. (ii) Electrolytes of type AB 2 (1:2 e.g., PbCl 2 , CaF2 Solubility product PbCl 2 ⇌ Pb 2   2Cl  x 2x In a saturated solution of sparingly soluble electrolyte two equilibria exist and can be represented as, AB ⇌ Ksp  [Pb 2  ][Cl  ]2 ; Ksp  [x ] [2 x ]2 ; Ksp  4 x 3 Solid AB Unionised (Dissolved ) ⇌ A  B ions [ A x By ]  K  ID CH 3 COOH ⇌ CH 3 COO  E3 suppress the dissociation of NH 4 OH due to common ion x  3 K sp / 4 type salt) Ionic Equilibrium (iii) Electrolyte of type A2B (2 : 1 type salt) e.g., Ag 2CrO4 , H 2 S Ag 2CrO4 ⇌ 2 Ag   CrO42  x 2x K sp Cl  The concentration of ions increases considerably in solution due to ionisation of HCl and due to common ion effect, dissociation of NaCl is 4 decreased. Hence, the ionic product [ Na  ][Cl  ] exceeds the solubility product of NaCl and therefore pure NaCl precipitates out from the solution. e.g., As 2 S 3 , Sb 2 S 3 As 2 S 3 ⇌ 2 As 3   3 S 2  2x 3x K sp  [ As 3  ]2 [S 2  ]3 ; K sp  [2 x ]2 [3 x ]3 ; K sp  4 x 2  27 x 3 ; x 5 RCOONa K sp Soap 108 e.g., AlCl 3 , Fe(OH )3 AlCl 3 ⇌ Al     3 Cl  (v) In qualitative analysis : The separation and identification of various basic radicals into different groups is based upon solubility product principle and common ion effect. ID 3x Ksp  [ Al 3 ][3Cl  ] ; Ksp  [x ] [3 x ]3 K sp 27. (a) Precipitation of group first radicals (Pb+2, Ag+ , (3) Criteria of precipitation of an electrolyte : When Ionic product of an electrolyte is greater than its solubility product, precipitation occurs. D YG (4) Applications of solubility product (i) In predicting the formation of a precipitate I : When Kip  K sp , then solution is unsaturated in which more solute can be dissolved. i.e., no precipitation. Case II : When Kip  K sp , then solution is saturated in which no more solute can be dissolved but no ppt. is fomed. K ip  K sp , then solution is U Case III : When Hg ) The group reagent is dilute HCl. [ Ag  ][Cl  ]  Ksp +2 U Ksp  27 x 4 ; x  4 ⇌ RCOO   Na  ; NaCl ⇌ Na   Cl  Hence, the ionic product [RCOO–] [Na+] exceeds the solubility product of soap and therefore, soap precipitates out from the solution. (v) Electrolyte of type AB 3 (1 : 3 type salt) x (iv) Salting out of soap : From the solution, soap is precipitated by the addition of concentrated solution of NaCl. E3 Ksp  108 x 5 60 (iv) Electrolyte of type A2 B3 (2 : 3 type salt) Case salt. HCl and NaCl dissociate into their respective ions as, NaCl ⇌ Na   Cl  ; HCl ⇌ H   Cl  Ksp  [ Ag  ]2 [CrO42  ] ; Ksp  [2 x ]2 [x ] ; Ksp  4 x 3 x 3 343 supersaturated and precipitation takes place. for AgCl. (b) Precipitation of group second radicals (Hg+2, Pb , Bi+3, Cu+2, Cd+2, As+3, Sb+3 and Sn+2) : The group reagent is in presence of dilute HCl. H 2S +2 [Pb 2 ][S 2 ]  Ksp for PbS. (c) Precipitation of group third radicals (Fe+3, Al+3 and Cr+3) The group reagent is NH 4 OH in presence of NH 4 Cl. [Fe 3 ][OH  ]3  Ksp (d) Precipitation of group fourth radicals (Co+2, Ni , Mn+2 and Zn+2) : The group reagent is H 2 S in +2 presence of NH 4 OH. (ii) In predicting the solubility of sparingly soluble salts Knowing the solubility product of a sparingly soluble salt at any given temperature, we can predict its solubility. (vi) Calculation of remaining concentration after precipitation : Sometimes an ion remains after precipitation if it is in excess. Remaining concentration can be determined, ST When the ionic product exceeds the solubility product, the equilibrium shifts towards left-hand side, i.e., increasing the concentration of undissociated molecules of the electrolyte. As the solvent can hold a fixed amount of electrolyte at a definite temperature, the excess of the electrolyte is thrown out from the solutions as precipitate. (iii) Purification of common salt : HCl gas is circulated through the saturated solution of common [Co 2 ][S 2 ]  Ksp (e) Precipitation of group fifth radicals (Ba+2, Sr+2, Ca ) The group reagent is ammonium carbonate in +2 presence of NH 4 Cl and NH 4 OH. [Ba 2 ] [CO32 ]  Ksp 344 Ionic Equilibrium [B  ] (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. ; K sp [Ca(OH ) 2 ] [OH  ] 2 In general [ A n  ]mleft  (ii) Limitations of Arrhenius concept K sp [ A m B n ] (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 m  ]n % precipitation of ion  Initial conc. - Remaining conc.  =   100 Initial conc.   (vii) Calculation of simultaneous solubility : Solubility of two electrolytes having common ion; when they are dissolved in the same solution, is called simultaneous solubility. Calculation of simultaneous solubility is divided into two cases. (2) Bronsted–Lowry concept : According to this concept, ID e.g., AgBr (Ksp  5  10 13 ) ; AgSCN (Ksp  10 12 ) “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.” Here, charge balancing concept is applied. Charge of Ag  = Charge of Br  + Charge of [ Ag  ] = [Br  ] (a  b ) = a U  [SCN  ] + HCl  H 2 O ⇌ H 3 O   Cl  b D YG e.g., CaF2 (Ksp  3.4  10 11 9 ) ; SrF2 (Ksp  2.9  10 ) Most of fluoride ions come of stronger electrolyte. Acid and Bases (1) Arrhenius concept : According to Arrhenius U 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. H   Cl  ; ST H2O (aq.) (aq.) NaOH (Base ) H2O Na   OH  (aq.) (aq ) 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, HNO 3 , H 2 SO 4 , HClO4 , etc., are examples NaOH , KOH , (CH 3 )4 NOH of are strong acids and strong bases. Every hydrogen compound cannot be regarded as an acid, e.g., CH 4 is not an acid. Similarly, CH 3 OH, C2 H5 OH , etc., have OH groups but they are not bases. …..(i) Base Acid Case II : When solubility products of two electrolytes are not close, i.e., they are not equally strong. HCl (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 equally strong (having close solubility product). ( Acid ) (b) The concept does not explain acidic and basic character of substances in non-aqueous solvents. certain salts such as AlCl 3 in aqueous solution. Case I : When the two electrolytes are almost SCN 60 [Ca 2  ]left  K sp [ AB ] E3 Example : [ A  ]left  CH 3 COOH  H 2 O ⇌ H 3 O   CH 3 COO  Acid …..(ii) Base (i) HCl and CH 3COOH are acids because they donate a proton to H 2O.(ii) NH 3 and CO 32  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 HCl – Cl– 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 OH–in 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, Ionic Equilibrium Base Acid Acid Base 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. (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 CO 2 , SO 2 , SO 3 etc and the basic oxides like 60 HCl  HF H 2Cl   F  (ii) It can explain the basic character of the substances like Na 2 CO 3 , NH 3 etc. CaO, BaO, MgO etc which take place even in the absence of the solvent e.g., CaO  SO 3 CaSO 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. E3 (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. 345 H 2 SO 4 (Sulphuric acid) H 3O  (Hydronium ion) HCN C6 H5 OH D YG (Hydrogen chloride) (Nitric acid) H 2O (Water) HSO 4 (Hydrogen sulphate ion) (Ortho phosphoric acid) H 3 PO 4 CH 3 COOH (Acetic acid) (Carbonic acid) H 2 CO 3 (Hydrogen sulphide) H2S NH 4 (Ammonium ion) U (Hydrogen cyanide) (Phenol) Conjugate base (Perchlorate ion) HSO 4 (Hydrogen sulphate ion)  Cl NO 3 (Chloride ion) (Nitrate ion) H 2O (Water) SO 42  (Sulphate ion) H 2 PO 4 (Dihydrogen phosphate ion) CH 3 COO  (Acetate ion) HCO 3 (Hydrogen carbonate ion) HS  NH 3 (Hydrogen sulphide ion) (Ammonia)  CN C6 H 5 O  (Cyanide ion) (Phenoxide ion) OH  (Hydroxide ion)  (Ethyl alcohol) C2 H 5 O NH 3 (Ammonia) NH 2 (Amide ion) (Methane) CH 3 (Methyl carbanion) ST C 2 H 5 OH CH 4 (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, H   OH  HOH Acid Base Increasing order of basic strength HCl HNO 3 ClO 4 U Acid (Perchloric acid) Increasing order of acidic strength HClO 4 ID Table: 9.1 Conjugate acid-base pairs (Ethoxide ion) 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 SO 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 BF3 , BCl 3 , AlCl 3 , BeCl 2 , etc. 346 Ionic Equilibrium (d) Molecules having a multiple bond between atoms of dissimilar electronegativity e.g., CO 2 , SO 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 : NH 3 ,  N H 2 , R  O  H.. (b) Negatively charged species or anions Acid strength of HA1  Acid strength of HA 2 K a1 K a2 Similarly, relative strengths of any two weak bases at the same concentration are given by the ratio of the square-roots of their dissociation constants. i.e., Basic strength of B OH 1  Basic strength of B O H 2 K b1 K b2 (1) Relative strength of Inorganic acids (i) Hydrides (a) The acidic strength increases with the increase in the electronegativity of the element directly attached with the hydrogen. H  F  H  OH  H  NH 2  H  CH 3 HCI  H 2 S  PH 3  SiH 4 (b) The acidic strength increases with the increase in atomic size, HF  HCl  HBr  HI ; H 2 O  H 2 S  H 2 Se  H 2 Te (ii) Oxyacids (a) Among oxyacids of the same type formed by different elements, acidic nature increases with increasing electronegativity, HOI  HOBr  HOCl ; HIO4  HBrO4  HClO 4 ID (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., Li  , Na  , Be 2  , Mg 2 , Al 3 BF3 , SO 3 etc.. On the contrary, a same concentration C , we have, 60 (c) Molecules in which the central atom has empty d-orbitals. e.g., SiF4 , SnCl 4 , PF5 etc. of ionisation constant K a1 and K a2 respectively at the E3 (b) All cations are expected to act as Lewis acids since they are deficient in electrons. U 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., Pb 2 , Cd 2 , Pt 2 , Hg 2 , Ro  , Rs  , I2 etc.. A Lewis base which holds its electrons strongly is called hard base, e.g., OH  , F  , H 2 O, NH 3 , CH 3 OCH 3 , etc. D YG 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 , CH 3 S  , (CH 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. U (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. ST (v) Limitations of lewis concept : It does not explain behaviour of well known protonic acids, as HCl, H 2 SO 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. Relative strength of acids and Bases In practice K a is used to define the strength only of those acids that are weaker than H 3 O  and K b is used to define the strength of only those bases that are weaker than OH . For two weak acids HA1 and HA 2 (b) In oxyacids of the same element, acidic nature increases with its oxidation number HOCl  HClO 2  HClO 3  HClO 4 ; H 2 SO 3  H 2 SO 4 1 3 5 7 HNO 2  HNO 3 (c) The strength of oxyacids increases from left to right across a period H 4 SiO 4  H 3 PO4  H 2 SO 4  HClO 4 (d) For the same oxidation state and configuration of the elements, acid strength decreases with increase in size of the atom. HNO3  HPO3 ; H3 PO4  H3 AsO4 HClO4  HBrO4  HIO4 (2) Relative strength of organic acids (i) A compound is acidic in nature, if its conjugate base can stabilize through resonance. Thus phenol is acidic while ethanol is neutral because the conjugate base of phenol (C 6 H 5 O  ) can be stabilized through resonance while that of alcohol (C 2 H 5 O  ) can not. (ii) Hydrogen atom attached to sp-hybridized carbon is more acidic than that on sp 2 hybridized carbon which in turn is more acidic than that on sp 3 hybridized carbon. Thus, HC  CH  CH 2  CH 2  CH 3  CH 3 sp sp 2 sp 3 (3) Relative strength of Inorganic bases Ionic Equilibrium (i) The basicity of a compound decreases with increase in electronegativity of the atom holding the...... (ii) The larger the size of the atom holding the unshared electrons, the lesser is the availability of electrons. (1) Simple salts : The salt formed by the interaction between acid and base, is termed as simple salt. These are of three types, (i) Normal salts : the salts formed by the loss of all possible protons (replaceable hydrogen atoms as H  ) are called normal salts. Such a salt does not contain either replacable hydrogen or a hydroxyl group. F   Cl   Br   I  ; O 2   S 2  Examples : NaCl , NaNO 3 , K 2 SO 4 , Ca 3 (PO 4 )2 , Na 3 BO 3 , (iii) Presence of negative charge on the atom holding the electron pair increases the basicity, while the presence of positive charge on the atom holding the electron pair decreases the basicity. Na 2 HPO 3 (one H atom is not replaceable as H 3 PO 2 is a dibasic acid) NaH 2 PO 2 (both H atoms are not 60.. electron pair, N H 3  H 2 O :  H F : 347 replaceable as H 3 PO 2 is a monobasic acid) etc. Be (OH )2  Mg(OH )2  Ca(OH )2  Sr (OH )2  Ba(OH )2 etc. CsOH is the strongest known base (iii) Basic salts : Salts formed by incomplete neutralisation of poly acidic bases are called basic salts. Such salts still contain one or more hydroxyl groups. These salts when neutralised by acids form normal salts. Examples: Zn(OH )Cl, Mg(OH )Cl, Fe(OH )2 Cl, Bi(OH )2 Cl ID (iv) Among alkali and alkaline earth hydroxides (oxides) the basic nature increases with electropositivity E3 LiOH  NaOH  KOH  RbOH  CsOH ; (ii) Acidic salts : Salts formed by incomplete neutralisation of poly-basic acids are called acidic salts. Such salts still contain one or more replaceable hydrogen atoms. These salts when neutralised by bases form normal salts. Examples : NaHCO 3 , NaHSO 4 , NaH 2 PO4 , Na 2 HPO 4 , OH   H 2 O  H 3 O  U (v) On going down the group; basic nature decreases with size of the central atom due to decrease in the ability to donate the lone pair. NH 3  PH 3  AsH 3  SbH 3  BiH 3 (4) Relative strength of Organic bases acid D YG (i) Higher the electron density on nitrogen, more is the basic character of amine. (ii) A compound is basic in nature, if its conjugate can be stabilized through resonance. Thus NH 2 | guanidine ( NH 3  C  NH ) is as strong alkali as metal NH 2  | hydroxides because its conjugate acid (H 3 N  C  NH ) (2) Double salts : The addition compounds formed by the combination of two simple salts are termed double salts. Such salts are stable in solid state only. Examples : Ferrous ammonium sulphate, Potash alum and other alums. (3) Complex salts : These are formed by combination of simple salts or molecular compounds. These are stable in solid state as well as in solutions. FeSO 4  6 KCN K 4 [Fe(CN )6 ] K 2 SO 4 Simple salts is very much stabilised through resonance. The acid-base neutralisation and Salt U The reaction between an acid and a base to form salt and water is termed neutralisation ST HCl(aq.)  NaOH (aq.) ⇌ NaCl (aq.) Sodium Chloride Salt  H 2 O(l) The process of neutralisation does not produce the resulting solution always neutral; no doubt it involves the interaction of H  and OH  ions. The nature of the resulting solution depends on the particular acid and the particular base involved in the reaction. Salts : Salts are regarded as compounds made up of positive and negative ions. The positive part comes from a base while negative part from an acid. Salts are ionic compounds.The salts can be classified into following classes, Complex salt (4) Mixed salts : The salt which furnishes more than one cation or more than one anion when dissolved in water is called a mixed salt. Na Na OCl \ SO ; NH Examples : Ca / ; PO 4 4 4 / \ K Cl H Ionic product of water Water is a weak electrolyte and undergoes selfionistion to a small extent. “The product of concentrations of H  and OH  ions in water at a particular temperature is known as ionic product of water.” It is designated as K w. H 2 O ⇌ H   OH  ; H  57.3 kJM 1 K [H  ][OH  ] ; K[H 2 O]  [H  ][OH  ] ; K w  [H  ][OH  ] [ H 2 O] 348 Ionic Equilibrium The value of K w increases with the increase of temperature, i.e., the concentration H+ and OH– ions increases with increase in temperature. The value of K w at 25 o C is 1  10 14 mole/litre. (i) pH values of the solutions do not give us immediate idea of the relative strengths of the solutions. A solution of pH =1 has a hydrogen ion concentration 100 times that of a solution pH = 3 (not Since pure water is neutral in nature, H  ion concentration must be equal to OH  ion concentration. [H  ]  [OH  ]  x or [H  ][OH  ]  x 2  1  10 14 or 2  10 5 N HCl   [H ]  [OH ] ; the solution is acidic [H  ]  [OH  ] ; the solution is basic solutions are 4.40 and 4.70 (not double). (ii) pH value zero is obtained in 1N solution of strong acid. If the concentration is 2 N, 3 N , 10 N , etc. the respective pH values will be negative. (iii) A solution of an acid having very low concentration, say 10 8 N , can not have pH 8, as shown by pH formula but the actual pH value will be less than 7. Hydrogen ion concentration – pH scale pK value : p stands for negative logarithm. Just as Sorensen, a Danish biochemist developed a scale 1 [H  ] pK a   log Ka and pK b   log Kb Weak acids have higher pK a values. Similarly U Just as pH indicates the hydrogen ion concentration, the pOH represents the hydroxyl ion concentration, i.e.,  D YG pOH   log[ OH ] Considering H  and OH  ion concentrations range over many negative powers of 10, it is convenient to express them as pH or pOH, the dissociation constant (K) values also range over many negative powers of 10 and it is convenient to write them as pK. Thus, pK is the negative logarithm of dissociation constant. ID to measure the acidity in terms of concentrations of H  in a solution. As defined by him, “pH of a solution is the negative logarithm to the base 10 of the concentration of H+ ions which it contains.” pH   log[ H  ] or pH  log the relationship, weak bases have higher pK b values For any conjugate acid-base pair in aqueous solution, Ka  Kb  Kw [H  ][OH  ]  Kw  1  10 14 pKa  pKb  pKw  14 (at 298o K) Taking log on both sides, we have   log[ H ]  log[ OH ]  log Kw  log(1  10 14 ) or  log[ H  ]  log[ OH  ]   log Kw   log(1  10 14 ) or pH  pOH  pK w  14 Neutral solution [H+] [OH–] pH pOH > 10–7 < 10–7 7 7 7 >7 10–7 pH of some materials 10 8 of the pH of 10 8 M HCl & M NaOH If we use the relation pH   log[ H 3 O  ] we get pH equal to 8, but this is not correct because an acidic solution connot have pH greater than 7. In this condition neglected. H  concentration of water cannot be    H water Therefore, [H  ]total  H Acid Since HCl is strong acid and completely ionised, Material pH Material pH Gastric juice 1.4 Rain water 6.5 Lemon juice 2.1 Pure water 7.0 Vinegar 2.9 Human saliva 7.0 Soft drinks 3.0 Blood plasma 7.4 Beer 4.5 Tears 7.4 Black coffee 5.0 Egg 7.8 Cow’s milk 6.5 Household ammonia 11.9 Limitations of pH scale Calculation [H  ]HCl  1  10 8 , [H  ]H 2O  10 7 [H  ]total  [H  ]HCl  [H  ]H2O  10 8  10 7  10 8 [1  10 ]  10 8  11 pH   log 10 8  log 11  6.958 Similarly if NaOH concentration is 10 8 M Then, [OH  ]total  [10 8 ]NaOH  [10 7 ]H2O [OH  ]  10 8  11 ; pOH  6. 96 pH  7. 04 Ionic Equilibrium Buffer solutions For acidic buffers, pH  pK a  log A solution whose pH is not altered to any great extent by the addition of small quantities of either sirong acid (H+ ions) or a sirong base (OH– ions) is called the buffer solution. It can also be defined as a solution of reserve acidity or alkalinity which resists change of pH upon the addition of small amount of 349 [salt] [acid ] When [salt]  10 , then, pH  1  pK a and [acid ] when [salt] 1 , then, pH  pK a  1  [acid ] 10 So weak acid may be used for preparing buffer solutions having pH values lying within the ranges (1) Types of buffer solutions : There are two types of buffer solutions, pK a  1 and pK a  1. The acetic acid has a pK a of about (i) Solutions of single substances : The solution of the salt of a weak acid and a weak base. Example : ammonium acetate 4.8; it may, therefore, be used for making buffer solutions with pH values lying roughly within the range 3.8 to 5.8. (CH 3 COONH 4 ) , For basic Buffers, pOH  pK b  log Example : CH 3 COOH  CH 3 COONa pH can be calculated by the application of formula, pH  pOH  14 pH of a buffer solution does not change with dilution but it varies with temperature because value of Kw changes with temperature. (5) Buffer capacity : The property of a buffer solution to resist alteration in its pH value is known as buffer capacity. It has been found that if the ratio U (b) Basic buffer : It is the solution of a mixture of a weak base and a salt of this weak base with a strong acid. Example : NH 4 OH  NH 4 Cl Knowing pOH , [salt] [base ] ID (a) Acidic buffer : It is the solution of a mixture of a weak acid and a salt of this weak acid with a strong base. E3 NH 4 CN act as a buffer. (ii) Solutions of Mixtures : These are further of two types, 60 acid or alkali. D YG (2) Buffer action : Buffer action is the mechanism by which added H+ ions or OH– ions are almost neutralised; so that pH practically remains constant. Reserved base of buffer neutralises the added H  ions while the reserved acid of buffer neutralises the added OH– ions. (3) Examples of buffer solutions (i) Phthalic acid + potassium hydrogen phthalate (ii) Citric acid + sodium citrate. U (iii) Boric acid + borax (sodium tetraborate). (iv) Carbonic acid (H 2 CO 3 ) + sodium hydrogen ST carbonate (NaHCO 3 ). This system is found in blood and helps in maintaining pH of the blood close to 7.4 ( pH value of human blood lies between 7.36 – 7.42; a change in pH by 0.2 units may cause death). (v) NaH 2 PO4  Na 3 PO4 (vi) NaH 2 PO 4  Na 2 HPO 4 (vii) Glycerine + HCl (viii) The pH value of gastric juice is maintained between 1.6 and 1.7 due to buffer system. (4) Henderson - Hasselbalch equation : pH of an acidic or a basic buffer can be calculated by Henderson- Hasselbalch equation. or [salt] [acid] [salt] is unity, the pH of a particular buffer does [base ] not change at all. Buffer capacity is defined quantitatively as number of moles of acid or base added in one litre of solution as to change the pH by unity, i.e., Buffer capacity ( )  Number of moles of acid or base added to 1 litre Change in pH Thus greater the buffer capacity, the greater is its capacity to resist change in pH value. Buffer capacity is greatest when the concentration of salt and weak acid/base are equal, or when pH  pK a or pOH  pK b. (6) Significance of buffer solutions (i) Buffer solutions are used for comparing colorimetrically the hydrogen ion concentration of unknown solutions. (ii) Acetic acid-sodium acetate is used in the removal of phosphate radical during the qualitative analysis of the mixture. (iii) NH 4 Cl / NH 4 OH buffer is used for the precipitation of hydroxides of third group of qualitative analysis. (iv) In industries, buffer solutions are used in the alcoholic fermentation (pH 5 to 6.5), tanning of leather, 350 Ionic Equilibrium Salt hydrolysis It is the reaction of the cation or the anion or both the ions of the salt with water to produce either acidic or basic solution. Hydrolysis is the reverse of neutralization. Neutr al Acidic solution Basic H O 1.2 – 2.8 Red Yellow Methyl yellow 2.9 – 4.0 Red Yellow No hydrolysis Methyl orange 3.1 – 4.4 Pink Yellow Cationic hydrolysis Anionic Methyl red 4.2 – 6.3 Red Yellow Litmus 5.0 – 8.0 Red Blue Bromothymol blue 6.0 – 7.6 Yellow Blue Phenol red 6.4 – 8.2 Yellow Red Thymol blue (base) 8.1 – 9.6 Yellow Blue Phenolphthalein 8.3 – 10.0 Colourless Pink Thymolphthalein 8.3 – 10.5 Colourless Blue Alizarin yellow R 10.1 – 12.0 Blue Yellow Nitramine 10.8 – 13.0 Colourless Orange, Brown base ID acid U D YG regarded as constant so, [HA][BOH ]  K[H 2 O]  K h [BA ] where K h is called the hydrolysis constant. U (2) Degree of hydrolysis : It is defined as the fraction (or percentage) of the total salt which is hydrolysed at equilibrium. For example, if 90% of a salt solution is hydrolysed, its degree of hydrolysis is 0.90 or as 90%. It is generally represented by ‘ h ’. Number of moles of the salt hydrolysed h Total number of moles of the salt taken ST Kw Ka Kb Exp. for pH 1 [log 2 Kw+log Ka– log C] K h  h  C      pH=– K h  h  C      pH=– h  (K h ) Base solution Thymol blue constant. Since water is present in very large excess in the aqueous solution, its concentration [H 2 O] may be Kh  Acid solution Yellow Applying the law of chemical equilibrium, we get [HA][BOH ] K, where K is the equilibrium [BA ][H 2 O] Kh=Kw / Kb Colour Red salt Exp. for h pH range 1.2 – 1.8 (1) Hydrolysis constant : The general equation for the hydrolysis of a salt (BA), BA  H 2 O ⇌ HA BOH Exp. for Kh Kh=Kw / Ka Indicator Cresol red hydrolysis Types of salt (i) Salt of weak acid and strong base (ii) Salt of strong acid and weak base (iii) Salt of weak acid and weak base An indicator is a substance, which is used to determine the end point in a titration. In acid-base titrations, organic substance (weak acids or weak bases) are generally used as indicators. They change their colour within a certain pH range. The colour change and the pH range of some common indicators are tabulated below Table : 9.2 Colour changes of indicators with pH E3 Salt 2 Aqueous Indicators 60 electroplating, manufacture of sugar, paper manufacturing etc., (v) In bacteriological research culture media are generally buffered to maintain the pH required for the growth of the bacteria being studied. (vi) In biological systems buffer system of carbonic acid and sodium bicarbonate is found in our blood. It maintains the pH of blood to a constant value (about 7.4) inspite of various acid and base-producing reactions going on in our body. 1 [log Kw– 2 log Kb+ log C] 1 [log Ka+ 2 log Kw– log Kb] pH=– (iv) Salts of strong acids and strong bases do not undergo hydrolysis (they undergo only ionization) hence the resulting aqueous solution is neutral. Two theories have been proposed to explain the change of colour of acid-base indicators with change in pH. (i) Ostwald’s Theory (ii) Quinonoid theory (1) Selection of suitable indicator or choice of indicator : In order to choose a suitable indicator, it is necessary to understand the pH changes in the titrations. The change in pH in the vicinity of the equivalence point is most important for this purpose. The curve obtained by plotting pH as ordinate against the volume of alkali added as abscissa is known as neutralisation or titration curve. The suitable indicators for the following titrations are, (i) Strong acid Vs strong base : Phenolphthalein (pH range 8.3 to 10.5), methyl red (pH range 4.4 – 6.5) and methyl orange (pH range 3.2 to 4.5). (ii) Weak acid Vs strong base : Phenolphthalein. (iii) Strong acid Vs weak base : Methyl red and methyl orange. (iv) Weak acid vs. weak base : No suitable indicator can be used for such a titration. Reason for use of different indicators for different systems : Indicators are either weak acids or weak bases and when dissolved in water their dissociated form acquires a colour different from that Ionic Equilibrium HIn(aq.) ⇌ H  (aq.)  In (aq.) Red Green  Let K In be the equilibrium constant K In  [H  ][In ] [HIn] or [HIn]   [In ] [H  ] K In  The human eye can detect the change in colour if the ratio of the two forms of indicator ranges between 0.1 to 10. [ HIn] [In ] [HIn] [In ]  1. 0 , the colour visible will be yellow  10 , the colour visible will be red.  0. 1 , the colour visible will be green. In other words, The colour visible will be red, when pH  pK In  1 The colour visible will be yellow, when pH  pK In  E3 [In ]  ID If, [HIn]  not 3 times. (ii) A 4 × 10–5 M HCl solution is twice concentrated as compared to 2 × 10–5 M HCl solution but pH values of these solutions are 4.4 and 4.7 respectively and not double. pH can be zero in 1M HCl or it can be negative for more concentrated solutions like 2M, 3M, 10M etc. At the temperature of the human body which is nearly 37°C, pH of neutral solution is 6.8. Buffer solutions have reserve acidity and reserve alkalinity. The greater the buffer capacity, the greater is its capacity to resist change in pH value. Buffers cannot withstand the addition of large amounts of acids or alkalies. The addition of 0.1mol per litre of [H+] or OH– is about the maximum that any buffer can be expected to withstand. 60 of the undissociated form. Consider a weak acid indicator of the general formula HIn, where in represents indicator. The equilibrium established in aqueous solution will be The colour visible will be green, when pH  pK In  1 Thus, our imaginary indicator will be red at any U pH which just falls below pK In  1 and green at any pH which just exceeds pK In  1. The indicator changes its D YG colour in the narrow pH range pK In  1 to pK In  1 from red to (red-yellow, yellow, yellow-green) green. We can therefore use this indicator to locate this narrow pH range. In other words, in order to use the indicator effectively in this range, we should have a solution for which pH is very near to pK In of the indicator. The colour change of an indicator can, therefore, be summarised as, Mid point of change KIn PKIn U [H+] pH First change of colour 10 KIn PKIn – 1 Colour change complete 0.1 KIn PKIn + 1 ST It is for this reason that we use different indicators for different systems.  pH of boiling water is 6.5625. It does not mean  that boiling water is not neutral. It is due to greater dissociation of H2O into H+ and OH–. pH values of solutions do not give the exact idea of their relative strengths e.g., (i) A solution with pH = 1 had [H+] 100 times than that with pH = 3 and 351