Current Affairs JEE Main & Advanced

"Electrochemical cell or Galvanic cell is a device in which a spontaneous redox reaction is used to convert chemical energy into electrical energy i.e. electricity can be obtained with the help of oxidation and reduction reaction". (1) Characteristics of electrochemical cell :  Following are the important characteristics of electrochemical cell, (i) Electrochemical cell consists of two vessels, two electrodes, two electrolytic solutions and a salt bridge. (ii) The two electrodes taken are made of different materials and usually set up in two separate vessels. (iii) The electrolytes are taken in the two different vessels called as half - cells. (iv) The two vessels are connected by a salt bridge/porous pot. (v) The electrode on which oxidation takes place is called the anode (or – ve pole) and the electrode on which reduction takes place is called the cathode (or + ve pole). (vi) In electrochemical cell, ions are discharged only on the cathode. (vii) Like electrolytic cell, in electrochemical cell, from outside the electrolytes electrons flow from anode to cathode and current flow from cathode to anode. (viii) For electrochemical cell, \[{{E}_{cell}}=+ve,\,\,\,\Delta G=-ve.\] (ix) In a electrochemical cell, cell reaction is exothermic. (2) Salt bridge and its significance (i) Salt bridge is U – shaped glass tube filled with a gelly like substance, agar – agar (plant gel) mixed with an electrolyte like KCl, KNO3, NH4NO3 etc. (ii) The electrolytes of the two half-cells should be inert and should not react chemically with each other. (iii) The cation as well as anion of the electrolyte should have same ionic mobility and almost same transport number, viz. \[KCl,\,KN{{O}_{3}},\,N{{H}_{4}}N{{O}_{3}}\]etc. (iv) The following are the functions of the salt bridge, (a) It connects the solutions of two half - cells and completes the cell circuit. (b) It prevent transference or diffusion of the solutions from one half cell to the other. (c) It keeps the solution of two half - cells electrically neutral. (d) It prevents liquid – liquid junction potential i.e. the potential difference which arises between two solutions when they contact with each other. (3) Representation of an electrochemical cell The cell may be written by arranging each of the pair left – right, anode – cathode, oxidation – reduction, negative and positive in the alphabetical order as, (4) Reversible and irreversible cells : A cell is said to be reversible if the following two conditions are fulfilled (i) The chemical reaction of the cell stops when an exactly equal external emf is applied. (ii) The chemical reaction of the cell is reversed and the current flows in opposite direction when the external emf is slightly higher than that of the cell. Any other cell, which does not obey the above two conditions, is termed as irreversible. Daniell cell is reversible but \[Zn\,|{{H}_{2}}S{{O}_{4}}|\,Ag\] cell is irreversible in nature (5) Types of electrochemical cells : Two main types more...

Auto Oxidation and Disproportionation   Autoxidation.             (1) Turpentine and numerous other olefinic compounds, phosphorus and certain metals like Zn and Pb can absorb oxygen from the air in presence of water. The water is oxidised to hydrogen peroxide. This phenomenon of formation of \[{{H}_{2}}{{O}_{2}}\] by the oxidation of \[{{H}_{2}}O\] is known as autoxidation. The substance such as turpentine or phosphorus or lead which can activate the oxygen is called activator. The activator is supposed to first combine with oxygen to form an addition compound, which acts as an autoxidator and reacts with water or some other acceptor so as to oxidise the latter. For example;                         \[\underset{(\text{activator})}{\mathop{Pb}}\,+{{O}_{2}}\to \,\underset{\text{(autoxidator)}}{\mathop{Pb{{O}_{2}}}}\,\]                   \[Pb{{O}_{2}}+\underset{(\text{acceptor})}{\mathop{{{H}_{2}}O}}\,\to PbO+{{H}_{2}}{{O}_{2}}\]             (2) The turpentine or other unsaturated compounds which act as activators are supposed to take up oxygen molecule at the double bond position to form unstable peroxide called moloxide, which then gives up the oxygen to water molecule or any other acceptor.                         \[\begin{matrix}    RCH=CHR+{{O}_{2}} & \to  & RHC & CHR  \\    {} & {} & \,\,\,\,\,\,\,O & O\,\,\,\,\,\,\,  \\ \end{matrix}\]                         \[\begin{array}{*{35}{l}}    RHC & CHR+2{{H}_{2}}O & \to  & RCH=CHR+2{{H}_{2}}{{O}_{2}}  \\    \,\,\,\,\,\,O & O & {} & {}  \\ \end{array}\]                                                 \[2KI+{{H}_{2}}{{O}_{2}}\to 2KOH+{{I}_{2}}\]             The evolution of iodine from KI solution in presence of turpentine can be confirmed with starch solution which turns blue.             (3) The concept of autoxidation help to explain the phenomenon of induced oxidation. \[N{{a}_{2}}S{{O}_{3}}\] solution is oxidised by air but \[N{{a}_{3}}As{{O}_{3}}\] solution is not oxidised by air. If a mixture of both is taken, it is observed both are oxidised. This is induced oxidation.                                                 \[N{{a}_{2}}S{{O}_{3}}+{{O}_{2}}\to N{{a}_{2}}S{{O}_{5}}\]                                                                               Moloxide                                         \[N{{a}_{2}}S{{O}_{5}}+N{{a}_{3}}As{{O}_{3}}\to N{{a}_{3}}As{{O}_{4}}+N{{a}_{2}}S{{O}_{4}}\]                                                                 \[N{{a}_{2}}S{{O}_{3}}+N{{a}_{3}}As{{O}_{3}}+{{O}_{2}}\to N{{a}_{2}}S{{O}_{4}}+N{{a}_{3}}As{{O}_{4}}\]   Disproportionation.             One and the same substance may act simultaneously as an oxidising agent and as a reducing agent with the result that a part of it gets oxidised to a higher state and rest of it is reduced to lower state of oxidation. Such a reaction, in which a substance undergoes simultaneous oxidation and reduction is called disproportionation and the substance is said to disproportionate.             Following are the some examples of disproportionation,     (1) \[{{H}_{2}}{{O}_{2}}+{{H}_{2}}{{\overset{-1}{\mathop{O}}\,}_{2}}={{H}_{2}}O+{{\overset{0}{\mathop{O}}\,}_{2}}\]   (2) \[4K\overset{+5}{\mathop{Cl}}\,{{O}_{3}}\to 3K\overset{+7}{\mathop{Cl}}\,{{O}_{4}}+\overset{-1}{\mathop{KCl}}\,\] (3) \[\overset{0}{\mathop{4P}}\,+3NaOH+3{{H}_{2}}O\to 3Na{{H}_{2}}{{\overset{+1\,\,\,}{\mathop{PO}}\,}_{2}}+\overset{-3\,\,\,\,\,\,\,\,\,}{\mathop{P{{H}_{3}}}}\,\] (4) \[\overset{0}{\mathop{3C{{l}_{2}}}}\,+\underset{(conc.)}{\mathop{6NaOH}}\,\xrightarrow{hot}5Na\overset{-1}{\mathop{Cl}}\,+Na\overset{+5\,\,\,\,\,}{\mathop{Cl{{O}_{3}}}}\,+3{{H}_{2}}O\]   Important applications of redox-reactions.             Many applications are based on redox reactions which are occuring in environment. Some important examples are listed below;             (1) Many metal oxides are reduced to metals by using suitable reducing agents. For example \[A{{l}_{2}}{{O}_{3}}\] is reduced to aluminium by cathodic reduction in electrolytic cell. \[F{{e}_{2}}{{O}_{3}}\] is reduced to iron in a blast furnace using coke.             (2)  Photosynthesis is used to convert carbon dioxide and water by chlorophyll of green plants in the presence of sunlight to carbohydrates.                                                 \[6C{{O}_{2(g)}}+6{{H}_{2}}{{O}_{(l)}}\,\,\xrightarrow{Chlorophyll}\,{{C}_{6}}{{H}_{12}}{{O}_{6(aq.)}}+6{{O}_{2}}_{(g)}\]             In this case, \[C{{O}_{2}}\] is reduced to carbohydrates and water is oxidised to oxygen. The light provides the energy required for the reaction.   (3) Oxidation of fuels is an important source of energy which satisfies our daily need of life.                                                             Fuels \[+{{O}_{2}}\to \,C{{O}_{2}}+{{H}_{2}}O+\]Energy more...

(1) Kohlrausch law states that, “At time infinite dilution, the molar conductivity of an electrolyte can be expressed as the sum of the contributions from its individual ions” i.e., \[\Lambda _{m}^{\infty }={{\nu }_{+}}\,\lambda _{+}^{\infty }+{{\nu }_{-}}\lambda _{-}^{\infty }\], where, \[{{\nu }_{+}}\] and \[{{\nu }_{-}}\] are the number of cations and anions per  formula unit of electrolyte respectively and,\[\lambda _{+}^{\infty }\] and \[\lambda _{-}^{\infty }\] are the molar conductivities of the cation and anion at infinite dilution respectively. The use of above equation in expressing the molar conductivity of an electrolyte is illustrated as, The molar conductivity of HCl at infinite dilution can be expressed as, \[\Lambda _{HCl}^{\infty }={{\nu }_{{{H}^{+}}}}\lambda _{{{H}^{+}}}^{\infty }+{{\nu }_{C{{l}^{-}}}}\lambda _{C{{l}^{-}}}^{\infty }\];  For HCl, \[{{\nu }_{{{H}^{+}}}}=1\] and \[{{\nu }_{C{{l}^{-}}}}=1.\] So, \[\Lambda _{HCl}^{\infty }=(1\times \lambda _{{{H}^{+}}}^{\infty })+(1\times \lambda _{C{{l}^{-}}}^{\infty })\];  Hence, \[\Lambda _{HCl}^{\infty }=\lambda _{{{H}^{+}}}^{\infty }+\lambda _{C{{l}^{-}}}^{\infty }\] (2) Applications of Kohlrausch's law : Some typical applications of the Kohlrausch's law are described  below, (i) Determination of \[\Lambda _{m}^{\infty }\] for weak electrolytes : The molar conductivity of a weak electrolyte at infinite dilution \[(\Lambda _{m}^{\infty })\] cannot be determined by extrapolation method. However, \[\Lambda _{m}^{\infty }\] values for weak electrolytes can be determined by using the Kohlrausch's equation. \[\Lambda _{C{{H}_{3}}COOH}^{\infty }=\Lambda _{C{{H}_{3}}COONa}^{\infty }+\Lambda _{HCl}^{\infty }-\Lambda _{NaCl}^{\infty }\] (ii) Determination of the degree of ionisation of a weak electrolyte : The Kohlrausch's law can be used for determining the degree of ionisation of a weak electrolyte at any concentration. If \[\lambda _{m}^{c}\] is the molar conductivity of a weak electrolyte at any concentration C and, \[\lambda _{m}^{\infty }\] is the molar conductivity of a electrolyte at infinite dilution. Then, the degree of ionisation is given by, \[{{\alpha }_{c}}=\frac{\Lambda _{m}^{c}}{\Lambda _{m}^{\infty }}=\frac{\Lambda _{m}^{c}}{({{\nu }_{+}}\lambda _{+}^{\infty }+{{\nu }_{-}}\lambda _{-}^{\infty })}\] Thus, knowing the value of \[\Lambda _{m}^{c}\], and \[\Lambda _{m}^{\infty }\] (From the Kohlrausch's equation), the degree of ionisation at any concentration \[({{\alpha }_{c}})\] can be determined. (iii) Determination of the ionisation constant of a weak electrolyte : Weak electrolytes in aqueous solutions ionise to a very small extent. The extent of ionisation is described in terms of the degree of ionisation \[(\alpha ).\]In solution, the ions are in dynamic equilibrium with the unionised molecules. Such an equilibrium can be described by a constant called ionisation constant. For example, for a weak electrolyte AB, the ionisation equilibrium is, \[AB\] ? \[{{A}^{+}}+{{B}^{-}}\]; If C is the initial concentration of the electrolyte AB in solution, then the equilibrium concentrations of various species in the solution are, \[[AB]=C(1-\alpha ),\] \[[{{A}^{+}}]=C\alpha \] and \[[{{B}^{-}}]=C\alpha \] Then, the ionisation constant of AB is given by, \[K=\frac{[{{A}^{+}}][{{B}^{-}}]}{[AB]}=\frac{C\alpha .C\alpha }{C(1-\alpha )}=\frac{C{{\alpha }^{2}}}{(1-\alpha )}\] We know, that at any concentration C, the degree of ionisation \[(\alpha )\] is given by, \[\alpha =\Lambda _{m}^{c}/\Lambda _{m}^{\infty }\] Then, \[K=\frac{C{{(\Lambda _{m}^{c}/\Lambda _{m}^{\infty })}^{2}}}{[1-(\Lambda _{m}^{c}/\Lambda _{m}^{\infty })]}=\frac{C{{(\Lambda _{m}^{c})}^{2}}}{\Lambda _{m}^{\infty }(\Lambda _{m}^{\infty }-\Lambda _{m}^{c})}\]; Thus, knowing \[\Lambda _{m}^{\infty }\] and \[\Lambda _{m}^{c}\] at any concentration, the ionisation constant (K) of the electrolyte can be determined. (iv) Determination of the solubility of a sparingly soluble more...

(1) Definition : “The fraction of the total current carried by an ion is known as transport number, transference number or Hittorf number may be denoted by sets symbols like t+ and t– or tc and ta or nc and na”. From this definition, \[{{t}_{a}}=\frac{\text{Current carried by an anion}}{\text{Total current passed through the solution}}\]  \[{{t}_{c}}=\frac{\text{Current carried by a cation}}{\text{Total current passed through the solution}}\] evidently, \[{{t}_{a}}+{{t}_{c}}=1.\] (2) Determination of transport number : Transport number can be determined by Hittorf's method, moving boundary method, emf method and  from ionic mobility. (3) Factors affecting transport number  A rise in temperature tends to bring the transport number of cation and anion more closer to 0.5 (4) Transport number and Ionic mobility : Ionic mobility or Ionic conductance is the conductivity of a solution containing 1 g ion, at infinite dilution, when two sufficiently large  electrodes are placed 1 cm apart. \[\text{Ionic mobilities }({{\lambda }_{a}}\,\text{or}\,{{\lambda }_{c}})\propto \,\text{speeds of ions}\,\text{(}{{u}_{a}}\text{or}\,{{u}_{c}})\] Unit of ionic mobility is Ohm–1 cm2 or V–1S-1cm2 Ionic mobility and transport number are related as,  \[{{\lambda }_{a}}\,\text{or}{{\lambda }_{c}}={{t}_{a}}\,\text{or}\,{{t}_{c}}\times {{\lambda }_{\infty }}\] Absolute ionic mobility is the mobility with which the ion moves under unit potential gradient. It's unit is \[cm\,{{\sec }^{-1}}\]. \[\text{Absolute ionic mobility }=\frac{\text{Ionic mobility}}{\text{96,500}}\]

Electricity is carried out through the solution of an electrolyte by migration of ions. Therefore, (1) Ions move toward oppositely charged electrodes at different speeds. (2) During electrolysis, ions are discharged or liberated in equivalent amounts at the two electrodes, no matter what their relative speed is. (3) Concentration of the electrolyte changes around the electrode due to difference in the speed of the ions. (4) Loss of concentration around any electrode is proportional to the speed of the ion that moves away from the electrode, so \[\frac{\text{Loss around anode}}{\text{Loss around cathode}}=\frac{\text{Speed of cation}}{\text{Speed of anion}}\]    The relation is valid only when the discharged ions do not react with atoms of the electrodes. But when the ions combine with the material of the electrode, the concentration around the electrode shows an increase.

In general, conductance of an electrolyte depends upon the following factors, (1) Nature of electrolyte : The conductance of  an electrolyte depends upon the number of ions present in the solution. Therefore, the greater the number of ions in the solution the greater is the conductance. The number of ions produced by an electrolyte depends upon its nature. The strong electrolytes dissociate almost completely into ions in solutions and, therefore, their solutions have high conductance. On the other hand, weak electrolytes, dissociate to only small extents and give lesser number of ions. Therefore, the solutions of weak electrolytes have low conductance.  (2) Concentration of the solution : The molar conductance of electrolytic solution varies with the concentration of the electrolyte. In general, the molar conductance of an electrolyte increases with decrease in concentration or increase in dilution. The molar conductance of strong electrolyte (\[HCl,KCl,KN{{O}_{3}})\] as well as weak electrolytes (\[C{{H}_{3}}COOH,N{{H}_{4}}OH)\] increase with decrease in concentration or increase in dilution. The variation is however different for strong and weak electrolytes. The variation of molar conductance with concentration can be explained on the basis of conducting ability of ions for weak and strong electrolytes. For weak electrolytes the variation of \[\Lambda \] with dilution can be explained on the bases of number of ions in solution. The number of ions furnished by an electrolyte in solution depends upon the degree of dissociation with dilution. With the increase in dilution, the degree of dissociation increases and as a result molar conductance increases. The limiting value of molar conductance \[({{\Lambda }^{0}})\] corresponds to degree of dissociation equal to 1 i.e., the  whole of the electrolyte dissociates. Thus, the degree of dissociation can be calculated at any concentration as, \[\alpha =\frac{{{\Lambda }^{c}}}{{{\Lambda }^{0}}}\] where \[\alpha \] is the degree of dissociation, \[{{\Lambda }^{c}}\] is the molar conductance at concentration C and \[{{\Lambda }^{0}}\] is the molar conductance at infinite dilution. For strong electrolytes, there is no increase in the number of ions with dilution because strong electrolytes are completely ionised in solution at all concentrations (By definition). However, in concentrated solutions of strong electrolytes there are strong forces of attraction between the ions of opposite charges called inter-ionic forces. Due to these inter-ionic forces the conducting ability of the ions is less in concentrated solutions. With dilution, the ions become far apart from one another and inter-ionic forces decrease. As a result, molar conductivity increases with dilution. When the concentration of the solution becomes very-very low, the inter-ionic attractions become negligible and the molar conductance approaches the limiting value called molar conductance at infinite dilution. This value is characteristic of each electrolyte. (3) Temperature : The conductivity of an electrolyte depends upon the temperature. With increase in temperature, the conductivity of an electrolyte increases.

When a voltage is applied to the electrodes dipped into an electrolytic solution, ions of the electrolyte move and, therefore, electric current flows through the electrolytic solution. The power of the electrolytes to conduct electric current is termed conductance or conductivity. (1) Ohm's law : This law states that the current flowing through a conductor is directly proportional to the potential difference across it, i.e., \[I\propto V\]    where I is the current strength (In Amperes) and V is the potential difference applied across the conductor (In Volts) or  \[I=\frac{V}{R}\] or \[V=IR\] where R is the constant of proportionality and is known as resistance of the conductor. It is expressed in Ohm's and is represented as \[\Omega .\]The above equation is known as Ohm's law. Ohm's law may also be stated as, “the strength of current flowing through a conductor is directly proportional to the potential difference applied across the conductor and inversely proportional to the resistance of the conductor.” (2) Resistance : It measures the obstruction to the flow of current. The resistance of any conductor is directly proportional to the length (l) and inversely proportional to the area of cross-section (a) so that \[R\propto \frac{l}{a}\,\,\,\,\text{or   }R=\rho \frac{l}{a}\] where \[\rho \](rho) is the constant of proportionality and is called specific resistance or resistivity. The resistance depends upon the nature of the material. Units : The unit of resistance is ohm \[(\Omega ).\]In terms of SI, base unit is equal to \[(kg{{m}^{2}})\,/\,({{s}^{3}}{{A}^{2}}).\] (3) Resistivity or specific resistance : We know that resistance R is \[R=\rho \frac{l}{a}\]; Now, if \[l=1\,cm,\,a=1\,c{{m}^{2}}\]then \[R=\rho \] Thus, resistivity is defined as the resistance of a conductor of 1 cm length and having area of cross-section equal to \[1\,c{{m}^{2}}.\] Units : The units of resistivity are \[\rho =R.\frac{a}{l}=Ohm\frac{c{{m}^{2}}}{cm}\] \[=Ohm.\,cm\] Its SI units are Ohm metre \[(\Omega \,m).\] But quite often Ohm centimetre \[(\Omega \,cm)\] is also used. (4) Conductance : It is a measure of the ease with which current flows through a conductor.  It is an additive property. It is expressed as G. It is reciprocal of the resistance, i.e., \[G=\frac{1}{R}\] Units : The units of conductance are reciprocal Ohm\[(oh{{m}^{-1}})\] or mho. Ohm is also abbreviated as \[\Omega \] so that \[Oh{{m}^{-1}}\] may be written as \[{{\Omega }^{-1}}.\] According to SI system, the units of electrical conductance is Siemens, S (i.e., \[1\text{S}=1\,{{\Omega }^{-1}}).\] (5) Conductivity : The inverse of resistivity is called conductivity (or specific conductance). It is represented by the symbol, \[\kappa \] (Greek kappa). The IUPAC has recommended the use of term conductivity over specific conductance. It may be defined as, the conductance of a solution of 1 cm length and having 1 sq. cm as the area of cross-section. In other words, conductivity is the conductance of one centimetre cube of a solution of an electrolyte. Thus, \[\kappa =\frac{1}{\rho }\]                                                               Units : The units of conductivity are \[\kappa =\frac{1}{Ohm.\,cm}=Oh{{m}^{-1}}\]cm–1  or \[{{\Omega }^{-1}}\,c{{m}^{-1}}\] In SI units, l is expressed in m area of cross-section in \[{{m}^{2}}\] so that the units of conductivity are more...