Current Affairs JEE Main & Advanced

(1) Postulates of Arrhenius theory (i) In aqueous solution, the molecules of an electrolyte undergo spontaneous dissociation to form positive and negative ions. (ii) \[\text{Degree of ionization (}\alpha \text{)}\] \[=\frac{\text{Number of dissociated molecules}}{\text{Total number of molecules of electrolyte before dissociation}}\] (iii) At moderate concentrations, there exists an equilibrium between the ions and undissociated molecules, such as,\[NaOH\] \[\rightleftharpoons \]  \[N{{a}^{+}}\] \[+\,O{{H}^{-}}\]; \[KCl\] \[\rightleftharpoons \]  \[{{K}^{+}}\] \[+\,C{{l}^{-}}\] 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 \[\alpha \] is nearly 1 for strong electrolytes, while it is very less than 1 for weak electrolytes. (ii) Higher the dielectric constant of a solvent more is its ionising power. Water is the most powerful ionising solvent as its dielectric constant is highest. (iii)  \[\alpha \ \text{ }\propto \frac{\text{1}}{\text{Con}\text{. of solution}}\]  \[\propto \frac{\text{1}}{\text{ wt}\text{. of solution}}\] \[\propto \] Dilution of solution \[\propto \] 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.

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 \[C{{H}_{3}}CO{{O}^{-}}\] ions and the following equilibrium is obtained, \[C{{H}_{3}}COOH+{{H}_{2}}O\]  \[\rightleftharpoons \]  \[C{{H}_{3}}CO{{O}^{-}}+{{H}_{3}}{{O}^{+}}\] Applying law of chemical equilibrium,  \[K=\frac{[C{{H}_{3}}CO{{O}^{-}}]\times [{{H}_{3}}{{O}^{+}}]}{[C{{H}_{3}}COOH]\times [{{H}_{2}}O]}\] 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, \[{{K}_{a}}=\frac{[C{{H}_{3}}CO{{O}^{-}}]\times [{{H}_{3}}{{O}^{+}}]}{[C{{H}_{3}}COOH]}\]                                                                                              …..(i) The fraction of total number of molecules of an electrolyte which ionise into ions is known as degree of dissociation/ionisation \[\alpha \]. If \['C'\] represents the initial concentration of the acid in moles \[{{L}^{-1}}\] and \[\alpha \] the degree of dissociation, then equilibrium concentration of the ions \[(C{{H}_{3}}CO{{O}^{-}}\] and \[{{H}_{3}}{{O}^{+}})\] is equal to \[C\alpha \] and that of the undissociated acetic acid \[=C(1-\alpha )\] i.e., we have                 \[C{{H}_{3}}COOH+{{H}_{2}}O\] \[\rightleftharpoons \]  \[C{{H}_{3}}CO{{O}^{-}}+{{H}_{3}}{{O}^{+}}\] Initial conc      \[C\]                                           0              0 Conc. at eqb. \[C(1-\alpha )\]                    \[C\alpha \]       \[C\alpha \] Substituting the values of the equilibrium concentrations in equation (i), we get \[{{K}_{a}}=\frac{C\alpha .C\alpha }{C(1-\alpha )}=\frac{{{C}^{2}}{{\alpha }^{2}}}{C(1-\alpha )}=\frac{C{{\alpha }^{2}}}{1-\alpha }\]                          …..(ii) In case of weak electrolytes, the value of \[\alpha \] is very small and can be neglected in comparison to 1 i.e., \[1-\alpha =1\]. Hence, we get \[{{K}_{a}}=C{{\alpha }^{2}}\] or \[\alpha =\sqrt{\frac{{{K}_{a}}}{C}}\]                                                                                                                   …..(iii) The degree of dissociation, \[\alpha \] can therefore be calcualted at a given concentration, \[C\] if \[{{K}_{a}}\] is known. Furher, if \[V\] is the volume of the solution in litres containing 1 mole of the electrolyte, \[C=1/V\]. Hence we have      \[\alpha =\sqrt{{{K}_{a}}V}\]                                                                                                                                                                                 …..(iv) Similarly, for a weak base like \[N{{H}_{4}}OH\], we have \[\alpha =\sqrt{{{K}_{b}}/C}=\sqrt{{{K}_{b}}V}\]                             more...

(1) Dissociation constant for weak acid : Consider an acid \[HA\] which, when dissolved in water ionizes as, \[HA\]  \[\rightleftharpoons \]  \[{{H}^{+}}+{{A}^{-}}\] Applying the law of mass action, \[{{K}_{a}}=\frac{[{{H}^{+}}][{{A}^{-}}]}{[HA]}\] 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. 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. \[{{H}_{3}}P{{O}_{4}}\]  \[\rightleftharpoons \] \[{{H}^{+}}+{{H}_{2}}PO_{4}^{-}\]              (I step) \[{{H}_{2}}PO_{4}^{-}\]  \[\rightleftharpoons \]  \[{{H}^{+}}+HPO_{4}^{-2}\]                      (II step) \[HPO_{4}^{-2}\]  \[\rightleftharpoons \] \[{{H}^{+}}+PO_{4}^{-3}\]                                   (III step) Let \[{{K}_{1}},\ {{K}_{2}}\] and \[{{K}_{3}}\] be the ionization constants of first, second and third steps respectively. Thus,         \[{{K}_{1}}=\frac{[{{H}^{+}}][{{H}_{2}}PO_{4}^{-}]}{[{{H}_{3}}P{{O}_{4}}]}\];\[{{K}_{2}}=\frac{[{{H}^{+}}][HPO_{4}^{-2}]}{[{{H}_{2}}PO_{4}^{-}]}\];\[{{K}_{3}}=\frac{[{{H}^{+}}][PO_{4}^{-3}]}{[HPO_{4}^{-2}]}\] In general, \[{{K}_{1}}>{{K}_{2}}>{{K}_{3}}\] The overall dissociation constant\[(K)\] is given by the relation, \[K={{K}_{1}}\times {{K}_{2}}\times {{K}_{3}}\] (2) Dissociation constant for weak base : The equilibrium of \[N{{H}_{4}}OH\] (a weak base) can be represented as, \[N{{H}_{4}}OH\]  \[\rightleftharpoons \] \[NH_{4}^{+}+O{{H}^{-}}\] Applying the law of mass action, \[{{K}_{b}}=\frac{[NH_{4}^{+}][O{{H}^{-}}]}{[N{{H}_{4}}OH]}\] \[{{K}_{b}}\] is constant at a definite temperature and does not change with the change of concentration.

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 \[N{{H}_{4}}Cl\] or \[NaOH\] to \[N{{H}_{4}}OH\] solution will suppress the dissociation of \[N{{H}_{4}}OH\] due to common ion either \[NH_{4}^{+}\] or \[O{{H}^{-}}\]. 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 \[O{{H}^{-}}\] ion concentration in third group.

If the concentration of the common ions in the solution of two electrolytes, for example \[{{H}^{+}}\] ion concentration in \[HCl\] and \[HN{{O}_{3}}\] or \[O{{H}^{-}}\] 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 \[H{{A}_{1}}\] and\[H{{A}_{2}}\]. Let \[{{V}_{1}}\] and \[{{V}_{2}}\] be their dilutions and \[{{\alpha }_{1}}\] and \[{{\alpha }_{2}}\] be their degree of dissociation at the respective dilution. Then, \[\frac{{{\alpha }_{1}}}{{{V}_{1}}}=\frac{{{\alpha }_{2}}}{{{V}_{2}}}\] Above equation is useful for calculating the relative dilution of two acids at which they would be isohydric.

In a saturated solution of sparingly soluble electrolyte two equilibria exist and can be represented as, \[\underset{\text{Solid}}{\mathop{AB}}\,\] ? \[\underset{\begin{smallmatrix}  \text{Unionised} \\  \text{(Dissolved)} \end{smallmatrix}}{\mathop{AB}}\,\]? \[\underset{\begin{smallmatrix}   \\  i\text{ons} \end{smallmatrix}}{\mathop{{{A}^{+}}+{{B}^{-}}}}\,\] Applying the law of mass action,\[\frac{[{{A}^{+}}][{{B}^{-}}]}{[AB]}=K\] 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,  \[[{{A}^{+}}][{{B}^{-}}]=K[AB]=K{K}'={{K}_{sp}}\] (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 \[{{A}_{x}}{{B}_{y}}\] which dissociates as, \[{{A}_{x}}{{B}_{y}}\]? \[x{{A}^{y+}}+y{{B}^{x-}}\] Applying law of mass action, \[\frac{{{[{{A}^{y+}}]}^{x}}{{[{{B}^{x-}}]}^{y}}}{[{{A}_{x}}{{B}_{y}}]}=K\] When the solution is saturated,  \[[{{A}_{x}}{{B}_{y}}]={K}'\] (constant) or \[{{[{{A}^{y+}}]}^{x}}{{[{{B}^{x-}}]}^{y}}=K[{{A}_{x}}{{B}_{y}}]=K{K}'={{K}_{sp}}\] (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., \[AgCl,\ BaS{{O}_{4}}\] \[AgCl\]? \[\underset{x}{\mathop{A{{g}^{+}}}}\,+\underset{x}{\mathop{C{{l}^{-}}}}\,\] \[{{K}_{sp}}=[A{{g}^{+}}][C{{l}^{-}}]\] ; \[{{K}_{sp}}={{x}^{2}}\]; \[x=\sqrt{{{K}_{sp}}}\]  (ii) Electrolytes of type\[A{{B}_{2}}\](1:2 type salt) e.g.,\[PbC{{l}_{2}},\ Ca{{F}_{2}}\] \[PbC{{l}_{2}}\] ? \[\underset{x}{\mathop{P{{b}^{2+}}}}\,+\underset{2x}{\mathop{2C{{l}^{-}}}}\,\] \[{{K}_{sp}}=[P{{b}^{2+}}]\,{{[C{{l}^{-}}]}^{2}}\];\[{{K}_{sp}}=[x]\ {{[2x]}^{2}}\];\[{{K}_{sp}}=4{{x}^{3}}\]  \[x=3\sqrt{{{K}_{sp}}/4}\] (iii) Electrolyte of type A2B (2 : 1 type salt) e.g.,\[A{{g}_{2}}Cr{{O}_{4}},\ {{H}_{2}}S\] \[A{{g}_{2}}Cr{{O}_{4}}\] ? \[\underset{2x}{\mathop{2A{{g}^{+}}}}\,+\underset{x}{\mathop{CrO_{4}^{2-}}}\,\] \[{{K}_{sp}}={{[A{{g}^{+}}]}^{2}}\ [CrO_{4}^{2-}]\];\[{{K}_{sp}}={{[2x]}^{2}}\ [x]\];\[{{K}_{sp}}=4{{x}^{3}}\]\[x=\sqrt[3]{\frac{{{K}_{sp}}}{4}}\] (iv) Electrolyte of type \[{{A}_{2}}{{B}_{3}}\](2 : 3 type salt)  e.g., \[A{{s}_{2}}{{S}_{3}},\ S{{b}_{2}}{{S}_{3}}\] \[A{{s}_{2}}{{S}_{3}}\]? \[\underset{2x}{\mathop{2A{{s}^{3+}}}}\,+\underset{3x}{\mathop{3{{S}^{2-}}}}\,\] \[{{K}_{sp}}={{[A{{s}^{3+}}]}^{2}}{{[{{S}^{2-}}]}^{3}}\] ; \[{{K}_{sp}}={{[2x]}^{2}}{{[3x]}^{3}}\]; \[{{K}_{sp}}=4{{x}^{2}}\times 27{{x}^{3}}\] \[{{K}_{sp}}=108{{x}^{5}}\] ; \[x=\sqrt[5]{\frac{{{K}_{sp}}}{108}}\] (v) Electrolyte of type \[A{{B}_{3}}\](1 : 3 type salt)  e.g.,\[AlC{{l}_{3}},\ Fe{{(OH)}_{3}}\] \[AlC{{l}_{3}}\]? \[\underset{x}{\mathop{A{{l}^{+++}}}}\,+\underset{3x}{\mathop{3C{{l}^{-}}}}\,\] \[{{K}_{sp}}=[A{{l}^{+3}}][3C{{l}^{-}}]\] ; \[{{K}_{sp}}=[x]\ {{[3x]}^{3}}\] \[{{K}_{sp}}=27{{x}^{4}}\]; \[x=\sqrt[4]{\frac{{{K}_{sp}}}{27}}\]. (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\[{{K}_{ip}}<{{K}_{sp}}\], then solution is unsaturated in which more solute can be dissolved. i.e., no precipitation. Case II : When \[{{K}_{ip}}={{K}_{sp}}\], then solution is saturated in which no more solute can be dissolved but no ppt. is fomed. Case III : When \[{{K}_{ip}}>{{K}_{sp}}\], 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. more...

(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. \[\underset{(Acid)}{\mathop{HCl}}\,\]        \[\underset{(aq.)}{\mathop{{{H}^{+}}}}\,+\underset{(aq.)}{\mathop{C{{l}^{-}}}}\,\] ;   \[\underset{(Base)}{\mathop{NaOH}}\,\]         \[\underset{(aq.)}{\mathop{N{{a}^{+}}}}\,+\underset{(aq)}{\mathop{O{{H}^{-}}}}\,\] 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. \[HCl,\ HN{{O}_{3}},\ {{H}_{2}}S{{O}_{4}},\,HCl{{O}_{4}}\], etc., are examples of strong acids and \[NaOH,\ KOH,\ {{(C{{H}_{3}})}_{4}}NOH\] are strong bases. Every hydrogen compound cannot be regarded as an acid, e.g., \[C{{H}_{4}}\] is not an acid. Similarly, \[C{{H}_{3}}OH,\ {{C}_{2}}{{H}_{5}}OH\], 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 \[AlC{{l}_{3}}\] in aqueous solution.   (2) Bronsted–Lowry concept : According to this concept, “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.” \[\underset{\text{Acid}}{\mathop{HC{{l}_{{}}}}}\,\,+\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\] ? \[{{H}_{3}}{{O}^{+}}+C{{l}^{-}}\]                                                                                                                                                               …..(i) \[\underset{\text{Acid}}{\mathop{C{{H}_{3}}COOH}}\,\,+\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\] ? \[{{H}_{3}}{{O}^{+}}+C{{H}_{3}}CO{{O}^{-}}\]                                                                                                                     …..(ii) (i) \[HCl\] and \[C{{H}_{3}}COOH\] are acids because they donate a proton to \[{{H}_{2}}O\].(ii) \[N{{H}_{3}}\] and \[CO_{3}^{2-}\] 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, (i) Protophilic more...

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 \[O{{H}^{-}}\].  For two weak acids \[H{{A}_{1}}\] and  \[H{{A}_{2}}\] of ionisation constant \[{{K}_{{{a}_{1}}}}\] and \[{{K}_{{{a}_{2}}}}\] respectively at the same concentration \[C\], we have, \[\frac{\text{Acid strength of }H{{A}_{\text{1}}}}{\text{Acid strength of }H{{A}_{2}}}=\sqrt{\frac{{{K}_{{{a}_{1}}}}}{{{K}_{{{a}_{2}}}}}}\] 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.,   \[\frac{\text{Basic strength of }B\text{O}{{\text{H}}_{\text{1}}}}{\text{Basic strength of }B\text{O}{{H}_{2}}}=\sqrt{\frac{{{K}_{{{b}_{1}}}}}{{{K}_{{{b}_{2}}}}}}\] (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. \[H-F>H-OH>H-N{{H}_{2}}>H-C{{H}_{3}}\]  \[HCI>{{H}_{2}}S>P{{H}_{3}}>Si{{H}_{4}}\] (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, \[HOI<HOBr<HOCl\];  \[HI{{O}_{4}}<HBr{{O}_{4}}<HCl{{O}_{4}}\] (b) In oxyacids of the same element, acidic nature increases with its oxidation number \[\underset{+1}{\mathop{HOCl}}\,<\underset{+3}{\mathop{HCl{{O}_{2}}}}\,<\underset{+5}{\mathop{HCl{{O}_{3}}}}\,<\underset{+7}{\mathop{HCl{{O}_{4}}}}\,\];\[{{H}_{2}}S{{O}_{3}}<{{H}_{2}}S{{O}_{4}}\] \[HN{{O}_{2}}<HN{{O}_{3}}\] (c) The strength of oxyacids increases from left to right across a period \[{{H}_{4}}Si{{O}_{4}}<{{H}_{3}}P{{O}_{4}}<{{H}_{2}}S{{O}_{4}}<HCl{{O}_{4}}\] (d) For the same oxidation state and configuration of the elements, acid strength decreases with increase in size of the atom. \[HN{{O}_{3}}>HP{{O}_{3}}\] ; \[{{H}_{3}}P{{O}_{4}}>{{H}_{3}}As{{O}_{4}}\] \[HCl{{O}_{4}}>HBr{{O}_{4}}>HI{{O}_{4}}\] (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 \[s{{p}^{2}}\] hybridized carbon which in turn is more acidic than that on \[s{{p}^{3}}\] hybridized carbon. Thus,  \[HC\equiv \underset{sp}{\mathop{CH}}\,>C{{H}_{2}}=\underset{s{{p}^{2}}}{\mathop{C{{H}_{2}}}}\,>C{{H}_{3}}-\underset{s{{p}^{3}}}{\mathop{C{{H}_{3}}}}\,\] (3) Relative strength of Inorganic bases (i) The basicity of a compound decreases with increase in electronegativity of the atom holding the electron pair, \[\overset{.\,\,\,.}{\mathop{N}}\,{{H}_{3}}>{{H}_{2}}\overset{.\,\,\,\,.}{\mathop{O}}\,:\ >\ H\underset{.\,\,\,\,.}{\overset{.\,\,\,\,.}{\mathop{F}}}\,:\] (ii) The larger the size of the atom holding the unshared electrons, the lesser is the availability of electrons. \[{{F}^{-}}>C{{l}^{-}}>B{{r}^{-}}>{{I}^{-}}\];  \[{{O}^{2-}}>{{S}^{2-}}\] (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. \[O{{H}^{-}}>{{H}_{2}}O>{{H}_{3}}{{O}^{+}}\] (iv) Among alkali and alkaline earth hydroxides (oxides) the basic nature increases with electropositivity \[LiOH<NaOH<KOH<RbOH<CsOH\]; \[Be{{(OH)}_{2}}<Mg{{(OH)}_{2}}<Ca{{(OH)}_{2}}<Sr{{(OH)}_{2}}<Ba{{(OH)}_{2}}\] \[CsOH\] 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.  \[N{{H}_{3}}>P{{H}_{3}}>As{{H}_{3}}>Sb{{H}_{3}}>Bi{{H}_{3}}\] (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 stabilized through resonance. Thus guanidine \[(N{{H}_{3}}-\overset{N{{H}_{2}}}{\mathop{\overset{|\ \ \ \ \ }{\mathop{C=}}\,}}\,NH)\] is as strong alkali as metal hydroxides because its conjugate acid \[({{H}_{3}}{{N}^{+}}-\overset{N{{H}_{2}}}{\mathop{\overset{|\ \ \ \ \ }{\mathop{C=}}\,}}\,NH)\] is very much stabilised through resonance.

The reaction between an acid and a base to form salt and water is termed neutralisation \[HC{{l}_{(aq.)}}+NaO{{H}_{(aq.)}}\] ? \[\underset{Salt}{\mathop{\underset{Sodium\,Chloride}{\mathop{NaC{{l}_{(aq.)}}}}\,}}\,+{{H}_{2}}{{O}_{(l)}}\] The process of neutralisation does not produce the resulting solution always neutral; no doubt it involves the interaction of \[{{H}^{+}}\] and \[O{{H}^{-}}\]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, (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. Examples :\[NaCl,\ NaN{{O}_{3}},{{K}_{2}}S{{O}_{4}},C{{a}_{3}}{{(P{{O}_{4}})}_{2}},N{{a}_{3}}B{{O}_{3}},\]  \[N{{a}_{2}}HP{{O}_{3}}\] (one \[H\] atom is not replaceable as \[{{H}_{3}}P{{O}_{2}}\] is a dibasic acid) \[Na{{H}_{2}}P{{O}_{2}}\] (both \[H\] atoms are not replaceable as \[{{H}_{3}}P{{O}_{2}}\] is a monobasic acid) etc. (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 : \[NaHC{{O}_{3}},NaHS{{O}_{4}},Na{{H}_{2}}P{{O}_{4}},N{{a}_{2}}HP{{O}_{4}}\], 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. \[\underset{\text{Simple}\ \text{salts}}{\mathop{FeS{{O}_{4}}+6KCN}}\,\to \underset{\text{Complex salt}}{\mathop{{{K}_{4}}[Fe{{(CN)}_{6}}]}}\,+{{K}_{2}}S{{O}_{4}}\] (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.  Examples : 

Water is a weak electrolyte and undergoes selfionistion to a small extent. “The product of concentrations of \[{{H}^{+}}\] and \[O{{H}^{-}}\] ions in water at a particular temperature is known as ionic product of water.” It is designated as \[{{K}_{w}}\]. \[{{H}_{2}}O\] ? \[{{H}^{+}}+O{{H}^{-}}\]; \[\Delta H=+57.3\ kJ{{M}^{-1}}\] \[K=\frac{[{{H}^{+}}][O{{H}^{-}}]}{[{{H}_{2}}O]}\];\[K[{{H}_{2}}O]=[{{H}^{+}}][O{{H}^{-}}]\];\[{{K}_{w}}=[{{H}^{+}}][O{{H}^{-}}]\] 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 \[1\times {{10}^{-14}}\]mole/litre. Since pure water is neutral in nature, \[{{H}^{+}}\] ion concentration must be equal to \[O{{H}^{-}}\] ion concentration. \[[{{H}^{+}}]=[O{{H}^{-}}]=x\] or \[[{{H}^{+}}][O{{H}^{-}}]={{x}^{2}}=1\times {{10}^{-14}}\] or \[x=1\times {{10}^{-7}}M\] or  \[[{{H}^{+}}]=[O{{H}^{-}}]=1\times {{10}^{-7}}\ mole\ litr{{e}^{-1}}\] 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, \[[{{H}^{+}}]=[O{{H}^{-}}]\]; the solution is neutral                      \[[{{H}^{+}}]>[O{{H}^{-}}]\]; the solution is acidic                        \[[{{H}^{+}}]<[O{{H}^{-}}]\]; the solution is basic