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60 E3 Chapter 9 Ionic Equilibrium (ii) Degree of ionization( ) 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 compound then it reacts with water t...
60 E3 Chapter 9 Ionic Equilibrium (ii) Degree of ionization( ) 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 compound then it reacts with water to form ions. Number of dissociated molecules Total number of molecules of electrolyte before dissociation (iii) At moderate concentrations, there exists an equilibrium between the ions and undissociated molecules, such as, NaOH ⇌ Na U Electrical conductors 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. OH ; KCl ⇌ K Cl This equilibrium state is called ionic equilibrium. (iv) Each ion behaves osmotically as a molecule. (2) Factors affecting degree of ionisation (2) Conductors which undergo decomposition (a chemical change) when an electric current is passed through them are known as electrolytic conductors or electrolytes. (i) At normal dilution, value of is nearly 1 for strong electrolytes, while it is very less than 1 for weak electrolytes. Electrolytes are further divided into two types on the basis of their strengths, (ii) Higher the dielectric constant of a solvent more is its ionising power. Water is the most powerful ionising solvent as its dielectric constant U (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. For example : HCl, H 2 SO 4 , NaCl, HNO3 , KOH , NaOH , ST HNO3 , AgNO3 , CuSO 4 etc. means all strong acids, bases and all types of salts. (ii) Substances which ionize to a small extent in their aqueous solution are known as weak electrolytes. Degree of ionization for this types of electrolytes is 1. For example : H 2O, CH 3COOH , NH 4 OH, HCN, Liq. SO 2 , HCOOH etc. means all weak acids and bases. 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. is highest. (iii) 1 1 Con. of solution wt. of solution Dilution of solution Amount of solvent (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 a strong electrolyte having a common ion. Ostwald's dilution law The strength of an acid or a bas is experimentally measured by determining its dissociation or ionisation constant. When acetic acid (a weak electrolyte) is dissolved in water, it dissociates partially into H or H 3 O and CH 3 COO ions and the following equilibrium is obtained, CH 3 COOH H 2O ⇌ CH 3 COO H 3 O Applying law of chemical equilibrium, K 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. [CH 3 COO ] [H 3 O ] [CH 3 COOH ] [H 2 O] In dilute solution, [H 2 O] is constant. The product of K and constant [H 2 O] is denoted as K a , the ionization constant or dissociation constant of the acid is, Ka [CH 3 COO ] [H 3 O ] [CH 3 COOH ] H 3 PO4 ⇌ H H 2 PO4 (I step) H 2 PO4 ⇌ H HPO42 (II step) HPO42 (III step) ⇌ H PO43 Let K1 , K 2 and K 3 be the ionization constants of first, second and third steps respectively. Thus, …..(i) [H ][PO43 ] [H ][HPO42 ] [H ][H 2 PO4 ] K ; K2 ; 3 [H 3 PO4 ] [HPO4 2 ] [H 2 PO4 ] K1 If ' C ' represents the initial concentration of the acid in moles L1 and the degree of dissociation, then equilibrium concentration of the In general, K1 K 2 K 3 The overall dissociation constant (K ) is given by the relation, ions (CH 3 COO and H 3 O ) is equal to C and that of the (2) Dissociation constant for weak base : The equilibrium of NH 4 OH (a weak base) can be represented as, CH 3 COOH H 2O ⇌ CH 3 COO H 3 O Initial conc 0 C K K1 K2 K3 E3 undissociated acetic acid C(1 ) i.e., we have 0 NH 4 OH ⇌ NH 4 OH Conc. at eqb. C(1 ) C C Substituting the values of the equilibrium concentrations in equation (i), we get is very small and can be …..(iii) solution in litres containing 1 mole of the electrolyte, C 1 / V. Hence we have …..(iv) Similarly, for a weak base like NH 4 OH , we have Kb / C Kb V …..(v) U The above equations lead to the following result ST “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, HA ⇌ H A Applying the law of mass action, K a Common ion effect 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 suppress the dissociation of NH 4 OH due to common ion either NH 4 or OH . CH 3 COOH ⇌ CH 3 COO H NH 4 OH ⇌ NH 4 OH CH 3 COONa CH 3 COO Na NH 4 Cl Common ion NH 4 Cl Common ion 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 adjust concentration of S 2 ions in second group and OH ion concentration in third group. Isohydric solution This is called Ostwald’s dilution law. K b is constant at a definite temperature and does not change with the change of concentration. U Ka C The degree of dissociation, can therefore be calcualted at a given concentration, C if K a is known. Furher, if V is the volume of the Ka V [ NH 4 ][OH ] [ NH 4 OH ] 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 D YG Ka C 2 or …..(ii) Applying the law of mass action, K b ID C .C C 2 2 C 2 C(1 ) C(1 ) 1 In case of weak electrolytes, the value of neglected in comparison to 1 i.e., 1 1. Hence, we get Ka 60 The fraction of total number of molecules of an electrolyte which ionise into ions is known as degree of dissociation/ionisation . [H ][ A ] [HA] 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. If the concentration of the common ions in the solution of two electrolytes , for example H ion concentration in HCl and HNO 3 or OH 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. 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 dilution. Then, 1 V1 2 K sp [ Ag ]2 [CrO42 ] ; Ksp [2 x ]2 [x ] ; K sp 4 x 3 V2 K sp Above equation is useful for calculating the relative dilution of two acids at which they would be isohydric. x 3 Solubility product (iv) Electrolyte of type A2 B3 (2 : 3 type salt) In a saturated solution of sparingly soluble electrolyte two equilibria exist and AB Unionised (Dissolved ) ⇌ A B As2 S 3 ⇌ 2 As 3 3 S 2 2x ions 3 2 K sp [ As ] [S [ A ][B ] K Applying the law of mass action, [ 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. K sp is termed as the solubility product. It is defined as the product [ A y ]x [B x ]y K [ A x By ] When the solution is saturated, [ A x By ] K (constant) or Ksp 108 x 5 ; x 5 K sp 108 (v) Electrolyte of type AB3 (1 : 3 type salt) AlCl3 ⇌ Al 3 Cl 3x x 3 Ksp [ Al ][3Cl ] ; Ksp [x ] [3 x ]3 K sp 27x 4 ; x 4 K sp. 27 (3) Criteria of precipitation of an electrolyte : When Ionic product of an electrolyte is greater than its solubility product, precipitation occurs. (4) Applications of solubility product (i) In predicting the formation of a precipitate Case I : When Kip Ksp , then solution is unsaturated in which U [ A y ]x [B x ]y K[ A x By ] KK K sp (constant) ] ; K sp [2 x ]2 [3 x ]3 ; K sp 4 x 2 27 x 3 ID of the concentration of ions in a saturated solution of an electrolyte at a given temperature. Consider, in general, the electrolyte of the type A x By which dissociates as, A x By ⇌ xA y yB x 3x 2 3 e.g., AlCl3 , Fe(OH )3 Hence, [ A ][B ] K[ AB] KK K sp (constant) Applying law of mass action, e.g., As2 S 3 , Sb 2 S 3 60 Solid E3 can be represented as, AB ⇌ 4 U D YG 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 ST AgCl ⇌ Ag Cl x x K sp [ Ag ][Cl ] ; K sp x 2 ; x K sp (ii) Electrolytes of type AB2 (1:2 type salt) e.g., PbCl2 , CaF2 PbCl2 ⇌ Pb2 2Cl x 2x K sp [Pb2 ][Cl ]2 ; Ksp [x ] [2 x ]2 ; K sp 4 x 3 x 3 K sp / 4 (iii) Electrolyte of type A B (2 : 1 type salt) e.g., Ag2CrO4 , H 2 S 2 Ag2CrO4 ⇌ 2 Ag CrO42 x 2x 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. Case III : When Kip K sp , then solution is supersaturated and precipitation takes place. 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. (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. (iii) Purification of common salt : HCl gas is circulated through the saturated solution of common salt. HCl and NaCl dissociate into their respective ions as, NaCl ⇌ Na Cl ; HCl ⇌ H Cl The concentration of Cl ions increases considerably in solution due to ionisation of HCl and due to common ion effect, dissociation of NaCl is decreased. Hence, the ionic product [ Na ][Cl ] exceeds the solubility product of NaCl and therefore pure NaCl precipitates out from the solution. (iv) Salting out of soap : From the solution, soap is precipitated by the addition of concentrated solution of NaCl. RCOONa ⇌ RCOO Na ; NaCl ⇌ Na Cl Soap Hence, the ionic product [RCOO ] [Na ] exceeds the solubility product of soap and therefore, soap precipitates out from the solution. – + (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. (a) Precipitation of group first radicals (Pb , Ag , Hg ) The group +2 + +2 reagent is dilute HCl. [ Ag ][Cl ] Ksp for AgCl. (b) Precipitation of group second radicals (Hg , Pb , Bi , Cu , Cd , As , Sb and Sn ) : The group reagent is H 2 S in presence of dilute HCl. +2 +2 +3 +2 +2 +3 +2 (c) Precipitation of group third radicals (Fe , Al and Cr ) The group reagent is NH 4 OH in presence of NH 4 Cl. +3 +3 +3 [Fe3 ][OH ]3 K sp (d) Precipitation of group fourth radicals (Co , Ni , Mn and Zn ) : The group reagent is H 2 S in presence of NH 4 OH. +2 +2 +2 +2 [Co 2 ][S 2 ] Ksp (e) Precipitation of group fifth radicals (Ba , Sr , Ca ) The group reagent is ammonium carbonate in presence of NH 4 Cl and NH 4 OH. +2 +2 +2 (vi) Calculation of remaining concentration after precipitation : Sometimes an ion remains after precipitation if it is in excess. Remaining concentration can be determined, [B ] NaOH Na OH H2O (aq.) (Base ) (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, HNO3 , H 2 SO 4 , HClO4 , etc., are examples of strong acids and NaOH, KOH , (CH 3 )4 NOH are strong bases. Every hydrogen ; [Ca 2 ]left K sp [Ca(OH ) 2 ] K sp [ Am Bn ] [B [OH ] 2 Similarly, CH 3 OH, C2 H 5 OH , etc., have OH groups but they are not bases. (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. (ii) Limitations of Arrhenius concept (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) The concept does not explain acidic and basic character of substances in non-aqueous solvents. (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 certain salts such as AlCl3 in aqueous solution. (2) Bronsted–Lowry concept : According to this concept, U K sp [ AB] “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.” + m n ] D YG In general (aq.) ID [Ba2 ] [CO 32 ] Ksp [ A n ]mleft (aq.) compound cannot be regarded as an acid, e.g., CH 4 is not an acid. [Pb2 ][S 2 ] K sp for PbS. Example : [ A ]left H Cl ; Initialconc. - Remaining conc. % precipitation of ion = 100 Initialconc. (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. Case I : When the two electrolytes are almost equally strong (having close solubility product). U e.g., AgBr (K sp 5 10 13 ) ; AgSCN (K sp 10 12 ) ST Charge of Ag = Charge of Br + Charge of SCN = [Br ] (a b) = a HCl H 2 O ⇌ H 3 O Cl …..(i) Base Acid CH 3 COOH H 2 O ⇌ H 3 O CH 3 COO Acid …..(ii) Base (i) HCl and CH 3 COOH are acids because they donate a proton to H 2 O.(ii) NH 3 and CO 32 are bases because they accept a proton from water. In reaction (i), in the reverse process, H O 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 H O – H O. These acid-base pairs are called conjugate acid-base pairs. + 3 – – + 3 2 Here, charge balancing concept is applied. [ Ag ] 60 +3 H2O E3 HCl (Acid) + [SCN ] 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 H O in aqueous solution fall to the strength of H O. 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, (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. + b Case II : When solubility products of two electrolytes are not close, i.e., they are not equally strong. e.g., CaF2 (K sp 3.4 10 11 ) ; SrF2 (K sp 2.9 10 9 ) Most of fluoride ions come of stronger electrolyte. Acid and Bases (1) Arrhenius concept : According to Arrhenius 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. + – + 3 3 – – 2 3 6 3 6 HCl HF H 2 Cl F Base Acid Acid Base (i) The protonic definition cannot be used to explain the reactions occuring in non-protonic solvents such as COCl , SO , N O , etc. (ii) It cannot explain the reactions between acidic oxides like CO 2 , SO 2 , SO 3 etc and the basic oxides like 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 BF , AlCl etc, do not have any hydrogen and hence cannot give a proton but are known to behave as acids. 2 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. (ii) It can explain the basic character of the substances like Na 2 CO 3 , NH 3 etc. (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 NH ). Limitations of Bronsted lowry concept Table: 9.1 Conjugate acid-base pairs Acid (Perchloric acid) HClO4 ClO4 3 2 (Hydrogen chloride) (Nitric acid) Cl NO 3 (Chloride ion) (Nitrate ion) H3O (Hydronium ion) H 2O (Water) HSO 4 (Hydrogen sulphate ion) SO 42 (Sulphate ion) H 3 PO4 (Ortho phosphoric acid) CH 3 COOH (Acetic acid) H 2 CO 3 (Carbonic acid) H2S (Hydrogen sulphide) NH 4 (Ammonium ion) HCN C6 H 5 OH (Hydrogen cyanide) (Phenol) NH 3 CH 4 ID H 2 PO4 (Dihydrogen phosphate ion) CH 3 COO (Acetate ion) HCO 3 (Hydrogen carbonate ion) HS NH 3 (Hydrogen sulphide ion) CN C6 H 5 O (Cyanide ion) (Phenoxide ion) OH (Hydroxide ion) U D YG C 2 H 5 OH E3 HCl HNO 3 Increasing order of acidic strength (Hydrogen sulphate ion) C2 H 5 O (Ammonia) (Ethoxide ion) (Ammonia) NH 2 (Amide ion) (Methane) CH 3 (Methyl carbanion) (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 U 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 Base ST Acid 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 , AlCl3 , BeCl 2 , etc. (b) All cations are expected to act as Lewis acids since they are deficient in electrons. (c) Molecules in which the central atom has empty dorbitals. e.g., SiF4 , SnCl 4 , PF5 etc. Increasing order of basic strength HSO 4 (Ethyl alcohol) 4 Conjugate base (Perchlorate ion) (Sulphuric acid) (Water) 2 3 H 2 SO 4 H 2O 2 60 HCl acts as acid in H O, stronger acid in NH , weak acid in CH COOH, neutral in C H and a weak base in HF. (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 (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 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., Pb2 , Cd 2 , Pt2 , Hg 2 , Ro , Rs , I2 etc.. A Lewis base which holds its electrons strongly is called hard base, e.g., OH , F , H 2O, NH 3 , CH 3 OCH 3 , etc. on the other hand, a Lewis 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 Kb is used to define the strength of only those bases that are weaker than OH . For two weak acids HA1 and HA 2 of ionisation constant K a1 and K a2 respectively at the same Acid strength of HA1 Acid strength of HA 2 D YG ST U HCI H 2 S PH3 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 HClO4 (b) In oxyacids of the same element, acidic nature increases with its oxidation number HOCl HClO2 HClO3 HClO4 ; H 2 SO 3 H 2 SO 4 5 sp 2 sp 3 (3) Relative strength of Inorganic bases (i) The basicity of a compound decreases with increase in electronegativity of the atom holding the electron pair,...... N H3 H2 O : H F :.. (ii) The larger the size of the atom holding the unshared electrons, the lesser is the availability of electrons. F Cl Br I ; O 2 S 2 (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. OH H 2 O H 3 O (iv) Among alkali and alkaline earth hydroxides (oxides) the basic nature increases with electropositivity CsOH is the strongest known base (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 BiH3 U 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 3 sp Be(OH)2 Mg(OH)2 Ca(OH)2 Sr(OH)2 Ba(OH)2 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., 1 Thus, HC CH CH 2 CH 2 CH 3 CH 3 LiOH NaOH KOH RbOH CsOH ; K a1 Basic strength of BOH 1 Basic strength of BOH 2 on sp 3 hybridized carbon. ID concentration C , we have, 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 60 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. (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. (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. conjugate base of phenol (C 6 H 5 O ) can be stabilized through resonance E3 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. 7 HNO 2 HNO 3 (c) The strength of oxyacids increases from left to right across a period H 4 SiO4 H 3 PO4 H 2 SO 4 HClO4 (d) For the same oxidation state and configuration of the elements, acid strength decreases with increase in size of the atom. HNO3 HPO3 ; H 3 PO4 H 3 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 (4) Relative strength of Organic bases (i) Higher the electron density on nitrogen, more is the basic character of amine. (ii) A compound is basic in nature, if its conjugate acid can be NH 2 | stabilized through resonance. Thus guanidine ( NH 3 C NH ) is as strong alkali as metal hydroxides because its conjugate acid NH 2 | (H 3 N C NH ) is very much stabilised through resonance. The acid-base neutralisation and Salt The reaction between an acid and a base to form salt and water is termed neutralisation 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, (1) Simple salts : The salt formed by the interaction between acid and base, is termed as simple salt. These are of three types, (replaceable hydrogen atoms as H ) are called normal salts. Such a salt does not contain either replacable hydrogen or a hydroxyl group. Examples : NaCl, NaNO 3 , K 2 SO 4 , Ca 3 (PO4 ) 2 , Na 3 BO 3 , Na 2 HPO3 (one H atom is not replaceable as H 3 PO2 is a dibasic Simple salts Complex salt D YG 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] U 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. Since pure + 1 [H ] Just as pH indicates the hydrogen ion concentration, the pOH represents the hydroxyl ion concentration, i.e., pH log[H ] or pH log pOH log[OH ] Considering the relationship, [H ][OH ] Kw 1 10 14 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 pKw 14 Acidic solution Neutral solution Basic solution Material Gastric juice Lemon juice Vinegar Soft drinks Beer Black coffee Cow’s milk U (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. OCl Na Na \ SO ; NH 4 PO4 Examples : Ca /\ ; 4 / Cl K H Ionic product of water + [H+] > 10–7 10–7 < 10–7 pH 1.4 2.1 2.9 3.0 4.5 5.0 6.5 ID acid) NaH 2 PO2 (both H atoms are not replaceable as H 3 PO2 is a monobasic acid) etc. (ii) Acidic salts : Salts formed by incomplete neutralisation of polybasic 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 HPO4 , etc. (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 (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 “pH of a solution is the negative logarithm to the base 10 of the concentration of H ions which it contains.” – ST 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 x 1 10 7 M or [H ] [OH ] 1 10 7 mole litre1 This shows that at 25 o C , in 1 litre only 10 7 mole of water is in ionic form out of a total of approximately 55.5 moles. Thus when, [H ] [OH ] ; the solution is neutral [H ] [OH ] ; the solution is acidic [H ] [OH ] ; the solution is basic 60 [OH–] pH < 10–7 10–7 >7 pH of some materials E3 (i) Normal salts : the salts formed by the loss of all possible protons Material Rain water Pure water Human saliva Blood plasma Tears Egg Household ammonia pOH >7 7
