Current Affairs JEE Main & Advanced

(1) Oxidation : Oxidation is a process which involves; addition of oxygen, removal of hydrogen, addition of non-metal, removal of metal, Increase in +ve valency, loss of electrons and increase in oxidation number. (i)  Addition of oxygen : 2Mg + O2 \[\to \]2MgO                     (ii) Removal of hydrogen : H2S+Cl2\[\to \]2HCl + S      (iii) Addition of Non-metal : Fe + S \[\to \] FeS               (iv) Removal of metal : 2KI+H2O2\[\to \]2KOH+I2 (v) Increase in +ve valency : \[F{{e}^{2+}}\,\to \,F{{e}^{3+}}+{{e}^{-}}\]                         (vi) Loss of electrons (also known as de-electronation) (a) \[{{H}^{0}}\to {{H}^{+}}+{{e}^{-}}\]            (Formation of proton) (b) \[MnO_{4}^{2-}\to MnO_{4}^{-}+{{e}^{-}}\]      (De-electronation of \[MnO_{4}^{2-}\]) (c) \[2F{{e}^{0}}\to 2Fe_{{}}^{3+}+6{{e}^{-}}\]      (De-electronation of iron) (vii) Increase in oxidation number (a) \[M{{g}^{0}}\to M{{g}^{2+}}\]                         (From 0 to +2) (b) \[{{\left[ F{{e}^{+2}}{{(CN)}_{6}} \right]}^{4-}}\to {{\left[ F{{e}^{+3}}{{(CN)}_{6}} \right]}^{3-}}\]   (From +2 to +3) (c) \[2C{{l}^{-}}\to Cl_{2}^{0}\]                             (From –1 to 0) (2) Reduction : Reduction is just reverse of oxidation. Reduction is a process which involves; removal of oxygen, addition of hydrogen, removal of non-metal, addition of metal, decrease in +ve valency, gain of electrons and decrease in oxidation number. (i) Removal of oxygen : \[CuO+C\to Cu+CO\]            (ii) Addition of hydrogen : \[C{{l}_{2}}+{{H}_{2}}\to 2HCl\] (iii) Removal of non-metal  \[2HgC{{l}_{2}}+SnC{{l}_{2}}\to H{{g}_{2}}C{{l}_{2}}+SnC{{l}_{4}}\] (iv) Addition of metal : \[HgC{{l}_{2}}+Hg\to H{{g}_{2}}C{{l}_{2}}\] (v) Decrease in +ve valency  (a) \[F{{e}^{3+}}\,\to \,F{{e}^{2+}}\]     (+ve valency decreases) (b) \[{{[Fe\,{{(CN)}_{6}}]}^{3-}}\to {{[Fe{{(CN)}_{6}}]}^{4-}}\] (–ve valency increases) (vi) Gain of electrons (also known as electronation) (a) \[Z{{n}^{2+}}(aq)+2{{e}^{-}}\to Zn(S)\] (Electronation of \[Z{{n}^{2+}}\]) (b) \[P{{b}^{2+}}+2{{e}^{-}}\to P{{b}^{0}}\] (Electronation of \[P{{b}^{2+}}\]) (c) \[{{[Fe{{(CN)}_{6}}]}^{3-}}+{{e}^{-}}\to {{[Fe{{(CN)}_{6}}]}^{4-}}\]                     (Electronation of \[{{[Fe{{(CN)}_{6}}]}^{3-}}\]) (vii) Decrease in oxidation number (a) \[M{{g}^{2+}}\to M{{g}^{0}}\]                         (From +2 to 0) (b) \[{{\left[ Fe{{(CN)}_{6}} \right]}^{3-}}\to {{\left[ Fe{{(CN)}_{6}} \right]}^{4-}}\]               (From +3 to +2) (c) \[Cl_{2}^{0}\to 2C{{l}^{-}}\]                             (From 0 to –1) (3) Redox-reactions (i) An overall reaction in which oxidation and reduction takes place simultaneously is called redox or oxidation-reduction reaction. These reactions involve transfer of electrons from one atom to another. Thus every redox reaction is made up of two half reactions; One half reaction represents the oxidation and the other half reaction represents the reduction. (ii) Types of redox reaction (a) Direct redox reaction : The reactions in which oxidation and reduction takes place in the same vessel are called direct redox reactions. (b) Indirect redox reaction : The reactions in which oxidation and reduction takes place in different vessels are called indirect redox reactions. Indirect redox reactions are the basis of electro-chemical cells. (c) Intermolecular redox reactions : In which one substance is oxidised while the other is reduced. For example, \[2\,Al+F{{e}_{2}}{{O}_{3}}\to A{{l}_{2}}{{O}_{3}}+2Fe\] Here, Al is oxidised to \[A{{l}_{2}}{{O}_{3}}\] while \[F{{e}_{2}}{{O}_{3}}\]is reduced to Fe. (d) Intramolecular redox reactions : In which one element of a compound is oxidised while the other is reduced. For example, \[2\,KCl{{O}_{3}}\xrightarrow{\Delta }2\,KCl+3\,{{O}_{2}}\] Here, \[C{{l}^{+5}}\] in \[KCl{{O}_{3}}\] is reduced to \[C{{l}^{-1}}\] in KCl while \[{{O}^{2-}}\] in \[KCl{{O}_{3}}\] is oxidised to \[O_{2}^{0}\].

(1) Molecular equations : When the reactants and products involved in a chemical change are written in molecular forms in the chemical equation, it is termed as molecular equation. Example :  \[Mn{{O}_{2}}+4\,HCl\to MnC{{l}_{2}}+2{{H}_{2}}O+C{{l}_{2}}\] In above example the reactants and products have been written in molecular forms, thus the equation is termed as molecular equation. (2) Ionic equations : When the reactants and products involved in a chemical change are ionic compounds, these will be present in the form of ions in the solution. The chemical change is written in ionic forms in chemical equation, it is termed as ionic equation. Example, \[Mn{{O}_{2}}+4\,{{H}^{+}}+4\,C{{l}^{-}}\]\[\to \]\[\,M{{n}^{2+}}+2C{{l}^{-}}+2{{H}_{2}}O+C{{l}_{2}}\] In above example the reactants and products have been written in ionic forms, thus the equation is termed as ionic equation. (3) Spectator ions : In ionic equations, the ions which do not undergo any change and equal in number in both reactants and products are termed as spectator ions and are not included in the final balanced equations. Example, \[Zn+2{{H}^{+}}+2C{{l}^{-}}\,\]\[\to \]\[Z{{n}^{2+}}+{{H}_{2}}+2C{{l}^{-}}\]    (Ionic equation) \[Zn+2{{H}^{+}}\] \[\to \]\[Z{{n}^{2+}}+{{H}_{2}}\]                           (Final ionic equation)            In above example, the \[C{{l}^{-}}\] ions are the spectator ions and hence are not included in the final ionic balanced equation.

(1) When metals are exposed to atmospheric conditions, they react with air or water in the environment to form undesirable compounds (usually oxides). This process is called corrosion. Almost all metals except the least active metals such as gold, platinum and palladium are attacked by environment i.e., undergo  corrosion. For example, silver tarnishes, copper develops a green coating, lead or stainless steel lose their lusture due to corrosion. Corrosion causes enormous damage to building, bridges, ships and many other articles made of iron. Thus corrosion is a process of deterioration of a metal as a result of its reaction with air or water (environment) surrounding it. In case of iron, corrosion is called rusting. Chemically, rust is hydrated form of ferric oxide, \[F{{e}_{2}}{{O}_{3}}\]. \[x{{H}_{2}}O\]. Rusting of iron is generally caused by moisture, carbon dioxide and oxygen present in air. It has been observed that rusting takes place only when iron is in contact with moist air. Iron does not rust in dry air and in vacuum. (2) Factors which affect corrosion : The main factors which affect corrosion are More the reactivity of metal, the more will be the possibility of the metal getting corroded. The impurities help in setting up voltaic cells, which increase the speed of corrosion Presence of electrolytes in water also increases the rate of corrosion Presence of \[C{{O}_{2}}\]in natural water increase rusting of iron. (v) When the iron surface is coated with layers of metals more active than iron, then the rate of corrosion is retarded. A rise in temperature (with in a reasonable limit) increases the rate of corrosion. (3) Classification of corrosion process : Depending upon the nature of corrosion, and the factors affecting it, the corrosion may be classified as follows. (i) Chemical corrosion : Such corrosion, generally takes place when (a) Reactive gases come in contact with metals at high temperatures e.g., corrosion in chemical industry. (b) Slow dissolution of metal takes place when kept in contact with non conducting media containing organic acids. (ii) Bio-chemical corrosion or Bio-corrosion: This is caused by the action of microorganisms. Soils of definite composition, stagnant water and certain organic products greatly favour the bio-corrosion. (iii) Electrochemical corrosion : It occurs in a gaseous atmosphere in the presence of moisture, in soils and in solutions. (4) Mechanism of rusting of iron : Electrochemical theory of rusting. The overall rusting involves the following steps, (i) Oxidation occurs at the anodes of each electrochemical cell. Therefore, at each anode neutral iron atoms are oxidised to ferrous ions. At anode : \[Fe(s)\xrightarrow{{}}F{{e}^{2+}}(aq)+2{{e}^{-}}.\] Thus, the metal atoms in the lattice pass into the solution as ions, leaving electrons on the metal itself. These electrons move towards the cathode region through the metal. (ii) At the cathodes of each cell, the electrons are taken up by hydrogen ions (reduction takes place). The \[{{H}^{+}}\] ions are obtained either from water or from acidic substances (e.g. \[C{{O}_{2}})\] in water \[{{H}_{2}}O\xrightarrow{{}}{{H}^{+}}+O{{H}^{-}}\]  more...

(1) The standard reduction potentials of a large number of electrodes have been measured using standard hydrogen electrode as the reference electrode. These various electrodes can be arranged in increasing or decreasing order of their reduction potentials. The arrangement of elements in order of increasing reduction potential values is called electrochemical series.It is also called activity series, of some typical electrodes. (2) Characteristics of Electrochemical series (i) The negative sign of standard reduction potential indicates that an electrode when joined with SHE acts as anode and oxidation occurs on this electrode. For example, standard reduction potential of zinc is –0.76 volt, When zinc electrode is joined with SHE, it acts as anode (–ve electrode) i.e., oxidation occurs on this electrode. Similarly, the +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode. (ii) The substances, which are stronger reducing agents than hydrogen are placed above hydrogen in the series and have negative values of standard reduction potentials. All those substances which have positive values of reduction potentials and placed below hydrogen in the series are weaker reducing agents than hydrogen. (iii) The substances, which are stronger oxidising agents than \[{{H}^{+}}\]ion are placed below hydrogen in the series. (iv) The metals on the top (having high negative value of standard reduction potentials) have the tendency to lose electrons readily. These are active metals. The activity of metals decreases from top to bottom. The non-metals on the bottom (having high positive values of standard reduction potentials) have the tendency to accept electrons readily. These are active non-metals. The activity of non-metals increases from top to bottom. Standard reduction electrode potentials at 298K   (3) Application of Electrochemical series (i) Reactivity of metals: The activity of the metal depends on its tendency to lose electron or electrons, i.e., tendency to form cation \[({{M}^{n+}})\]. This tendency depends on the magnitude of standard reduction potential. The metal which has high negative value  (or smaller positive value) of standard reduction potential readily loses the electron or electrons and is converted into cation. Such a metal is said to be chemically active. The chemical reactivity of metals decreases from top to bottom in the series. The metal higher in the series is more active than the metal lower in the series. For example, (a) Alkali metals and alkaline earth metals having high negative values of standard reduction potentials are chemically active. These react with cold water and evolve hydrogen. These readily dissolve in acids forming corresponding salts and combine with those substances which accept electrons. (b) Metals like Fe, Pb, Sn, Ni, Co, etc., which lie a little down in the series do not react with cold water but react with steam to evolve hydrogen.     (c)  Metals like Cu, Ag and Au  which lie below hydrogen are less reactive and do not  evolve hydrogen from water. (ii) Electropositive character of more...

The electrical work (electrical energy) is equal to the product of the EMF of the cell and electrical charge that flows through the external circuit i.e., \[{{W}_{\max }}=nF{{E}_{cell}}\]                                                                ......(i) According to thermodynamics the free energy change \[(\Delta G)\] is equal to the maximum work. In the cell work is done on the surroundings by which electrical energy flows through the external circuit, So \[-{{W}_{\max ,}}=\Delta G\]                                                                     ......(ii) from eq. (i) and (ii) \[\Delta G=-nFE_{cell}^{{}}\] In standard conditions  \[\Delta {{G}^{0}}=-\,nFE_{cell}^{0}\] Where \[\Delta {{G}^{0}}=\]standard free energy change But \[E_{cell}^{0}=\frac{2.303}{nF}RT\,\log {{K}_{c}}\] \[\therefore \]\[\Delta {{G}^{0}}=-nF\times \frac{2.303}{nF}RT\,\log \,{{K}_{c}}\] \[\Delta {{G}^{0}}=-\text{ 2}\text{.303 RT log }{{\text{K}}_{\text{c}}}\text{ }\]or \[\Delta G=\Delta G{}^\circ +2.303RT\log Q\] \[\Delta {{G}^{0}}=-RT\,\ln \,{{K}_{c}}\,\,\,\,\,\,\,\,(2.303\,\log X=\ln \,X)\]

(1) Nernst’s equation for electrode potential The potential of the electrode at which the reaction, \[{{M}^{n+}}(aq)+n{{e}^{-}}\to M(s)\] takes place is described by the equation,  \[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{RT}{nF}\ln \frac{[M(s)]}{[{{M}^{n+}}(aq.)]}\] or\[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{2.303\,\,RT}{nF}\log \frac{[M(s)]}{[{{M}^{n+}}(aq)]}\] above eq. is called the Nernst equation. Where, \[{{E}_{{{M}^{n+}}/M}}\]= the potential of the electrode at a given concentration, \[E_{{{M}^{n+}}/M}^{0}\] = the standard electrode potential R =  the universal gas constant, \[8.31\ J\,{{K}^{-1}}\,mo{{l}^{-1}}\] T= the temperature on the absolute scale, n = the number of electrons involved in the electrode reaction, F = the Faraday constant : (96500 C), \[[M(s)]\]= the concentration of the deposited metal, \[[{{M}^{n+}}(aq)]\]= the molar concentration of the metal ion in the solution, The concentration of pure metal M(s) is taken as unity. So, the Nernst equation for the \[{{M}^{n+}}/M\] electrode is written as,  \[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{2.303\,\,RT}{nF}\log \frac{1}{[{{M}^{n+}}(aq)]}\] At 298 K, the Nernst equation for the  \[{{M}^{n+}}/M\] electrode can be written as, \[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{0.0591}{n}\log \frac{1}{[{{M}^{n+}}(aq)]}\] For an electrode (half - cell) corresponding to the electrode reaction, Oxidised form \[+n{{e}^{-}}\to \]Reduced form The Nernst equation for the electrode is written as, \[{{E}_{half-cell}}=E_{half-cell}^{0}-\frac{2.303\,RT}{nF}\log \frac{[\text{Reduced form }]}{\text{ }\!\![\!\!\text{ Oxidised form }\!\!]\!\!\text{ }}\] At 298 K,  the Nernst equation can be written as, \[{{E}_{half-cell}}=E_{half-cell}^{0}-\frac{0.0591}{n}\log \frac{[\text{Reduced form }]}{\text{ }\!\![\!\!\text{ Oxidised form }\!\!]\!\!\text{ }}\] (2) Nernst’s equation for cell EMF For a cell in which the net cell reaction involving n electrons is, \[aA+bB\to cC+dD\] The Nernst equation is written as, \[{{E}_{cell}}=E_{cell}^{0}-\frac{RT}{nF}\text{ln}\frac{{{\text{ }\!\![\!\!\text{ C }\!\!]\!\!\text{ }}^{\text{c}}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}\] Where, \[E_{cell}^{0}=E_{cathode}^{0}-E_{anode}^{0}\].  The \[E_{cell}^{o}\] is called the standard cell potential. or  \[{{E}_{\text{cell}}}=E_{cell}^{o}-\frac{2.303\,RT}{nF}\log \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}\] At 298 K, above eq. can be written as, or \[{{E}_{\text{cell}}}=E_{cell}^{o}-\frac{0.0592}{n}\log \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}\] It may be noted here, that the concentrations of A, B, C and D referred in the eqs. are the concentrations at the time the cell emf is measured. (3) Nernst’s equation for Daniells cell :  Daniell’s cell consists of zinc and copper electrodes. The electrode reactions in Daniell’s cell are, At anode :                    \[Zn(s)\to Z{{n}^{2+}}(aq)+2{{e}^{-}}\] At cathode : \[C{{u}^{2+}}(aq)+2{{e}^{-}}\to Cu(s)\] Net cell reaction :        \[Zn(s)+C{{u}^{2+}}(aq)\to Cu(s)+Z{{n}^{2+}}(aq)\] Therefore, the Nernst equation for the Daniell’s cell is, \[{{E}_{cdll}}=E_{cell}^{0}-\frac{2.303\,RT}{2F}\log \frac{[Cu(s)][Z{{n}^{2+}}(aq)]}{[Zn(s)][C{{u}^{2+}}(aq)]}\] Since, the activities of pure copper and zinc metals are taken as unity, hence the Nernst equation for the Daniell’s cell is, \[{{E}_{cdll}}=E_{cell}^{0}-\frac{2.303\,RT}{2F}\log \frac{[Z{{n}^{2+}}(aq]}{[C{{u}^{2+}}(aq)]}\] The above eq. at 298 K is, \[{{E}_{cdll}}=E_{cell}^{o}-\frac{0.0591}{2}\log \frac{[Z{{n}^{2+}}(aq]}{[C{{u}^{2+}}(aq)]}V\] For Daniells cell,  \[E_{cell}^{0}=1.1\,V\] (4) Nernst's equation and equilibrium constant        For a cell, in which the net cell reaction involving n electrons is,  \[aA+bB\to cC+dD\]                                                           The Nernst equation is \[{{E}_{Cell}}=E_{cell}^{0}-\frac{RT}{nF}\ln \,\frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}\]                                        .....(i) At equilibrium, the cell cannot perform any useful work. So at equilibrium, \[{{E}_{Cell}}\]is zero. Also at equilibrium, the ratio \[\frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}={{\left[ \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}} \right]}_{equilibrium}}={{K}_{c}}\]

(1) “The difference in potentials of the two half – cells of a cell known as electromotive force (emf) of the cell or cell potential.” The difference in potentials of the two half – cells of a cell arises due to the flow of electrons from anode to cathode and flow of current from cathode to anode. \[Anode\underset{Flow\,\,of\,\,current}{\overset{Flow\,\,of\,\,electrons}{\longleftrightarrow}}Cathode\] (2) The emf of the cell or cell potential can be calculated from the values of electrode potentials of two the half – cells constituting the cell. The following three methods are in use : (i) When oxidation potential of anode and reduction potential of cathode are taken into account \[E_{\text{cell}}^{0}=\] Oxidation potential of anode + Reduction potential of cathode \[=E_{\text{ox}}^{0}\,(\text{anode})+E_{\text{red}}^{\text{0}}(\text{cathode})\] (ii) When reduction potentials of both electrodes are taken into account \[=E_{\text{Cathode}}^{\text{0}}-E_{\text{Anode}}^{\text{0}}\]\[=E_{\text{right}}^{\text{o}}-E_{\text{left}}^{o}\] (iii) When oxidation potentials of both electrodes are taken into account \[E_{\text{cell}}^{o}=\] Oxidation potential of anode – Oxidation potential of cathode \[=E_{\text{ox}}^{0}(\text{anode})-E_{\text{ox}}^{0}(\text{cathode})\]       (3) Difference between emf and potential difference

Emf	Potential difference
It is the potential difference between two electrodes when no current is flowing in the circuit.	It is the difference of the electrode potentials of the two electrodes when the cell is under operation.
It is the maximum voltage that the cell can deliver.	It is always less then the maximum value of voltage which the cell can deliver.
It is responsible for the steady flow of current in the cell.	It is not responsible for the steady flow of current in the cell.

(4) Cell EMF and the spontaneity of the reaction :         We know, \[\Delta G=-nF{{E}_{cell}}\]

more... Catgeory: JEE Main & Advanced Current Affairs Categories National International Agreements Appointments Government Schemes Summits Awards & Honours Business Economy & Banking Defense Sports Ranks & Reports Important Days Persons Science & Technology State Environment Bills & Acts Schemes Miscellaneous Archive August 2020(12) February 2020(215) January 2020(211) July 2020(35) June 2020(143) March 2020(99) May 2020(49) April 2019(263) August 2019(324) December 2019(331) February 2019(280) January 2019(306) July 2019(359) June 2019(231) March 2019(259) May 2019(258) November 2019(228) October 2019(222) September 2019(322) April 2018(298) August 2018(290) December 2018(292) July 2018(306) June 2018(311) March 2018(5) May 2018(328) November 2018(277) October 2018(276) September 2018(281) Trending Current Affairs Smart Cities India 2018 Expo Underway In New Delhi CII Pegs India's GDP Growth At 7.3-7.7% For 2018-19 World Environment Day 2018 Celebrated In New Delhi President 3-Nation Visit: India & Cuba Sign MoUs SCM BrahMos Test Fired Under Extreme Conditions 4 Indian-Origin Names In List Of Fortune's 40 Young Influential People Colombia joins NATO as global partner Bangladesh Cabinet Clears Deal With India For Port Usage PM Inaugurates Key Development Projects In Varanasi Paisabazaar.com Launches India’s First ‘Chance of Approval’ Feature Please Wait you are being redirected.... You need to login to perform this action. You will be redirected in 3 sec × Reset Password. Email : Send Password Cancel Add Mobile Enter Mobile Number Submit Logout Enter OTP OTP has been sent to your mobile number and is valid for one hour Enter One Time Password Submit Change Mobile: Submit Logout Mobile Number Verified Your mobile number is verified. Close Free Videos × For free video lectures fill in the form below and we will get back to you. * Email Address: * Mobile Number: Submit Cancel