How Can An Electrochemical Series Be Established Biology Essay

By mensurating the potencies of assorted electrodes versus stand ard H electrode ( SHE ) , a series of standard electrode potencies has been established. When the electrodes ( metals and non-metals ) in contact with their ions are arranged oh the footing of the values of their standard decrease potencies or standard oxidization potencies, the ensuing series is called theA electrochemical orA electromotiveA orA activity seriesA of the elements.

By international convention, the standard potencies of electrodes are tabulated for decrease half reactions, bespeaking the inclinations of the electrodes to act as cathodes towards HER. Those with positive EA° values for decrease half reactions do in fact act as cathodes versus SHE, while those with negative EA° values of decrease half reactions behave alternatively as anodes versus SHE. The electrochemical series is shown in the follow ing tabular array.

Standard Aqueous Electrode Potentials at 25A°C ‘The Electrochemical Series ‘

Component

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Electrode Reaction

( Decrease )

Standard Electrode Reduction potency

Eo, V

Lithium

K

Calcium

Sodium

Milligram

Aluminum

Zinc

Chromium

Fe

Cadmium

Nickel

Tin

H2

Copper

I2

Silver

Mercury

Br2

Cl2

Gold

F2

Li+A + e-A = Li

K+A + e-A = K

Ca2+A + 2e-A = Ca

Na+A + e-A = Na

Mg2+A + 2e-A = Mg

Al3+A + 3e-A = Al

Zn2+A + 2e-A = Zn

Cr3+A + 3e-A = Cr

Fe2+A + 2e-A = Fe

Cd2+A + 2e-A = Cadmium

Ni2+A + 2e-A = Ni

Sn2+A + 2e-A = Sn

2H+A + 2e-A = H2

Cu2+A + 2e-A = Cu

I2A + 2e-A = 2I-

Ag+A + e-A = Ag

Hg2+A + 2e-A = Hg

Br2A + 2e-A = 2Br-

Cl2A + 2e-A = 2Cl-

Au3+A + 3e-A = Au

F2A + 2e-A = 2F-

-3.05

-2.925

-2.87

-2.714

-2.37

-1.66

-0.7628

-0.74

-0.44

-0.403

-0.25

-0.14

0.00

+0.337

+0.535

+0.799

+0.885

+1.08

+1.36

+1.50

+2.87

A

Features of Electrochemical series

( ! ) The negative mark of standard decrease potency indicates that an electrode when joined with SHE acts as anode and oxidization occurs on this electrode. For illustration, standard decrease potency of Zn is -0.76 V. When Zn electrode is joined with SHE, it acts as anode ( -ve electrode ) i.e. , oxidization occurs on this electrode. Similarly, the +ve mark of standard decrease potency indicates that the electrode when joined with SHE acts as cathode and decrease occurs on this electrode.

( two ) A The substances which are stronger cut downing agents than H are placed above H in the series and have negative values of standard decrease potencies. All those substances which have positive values of decrease potencies and placed below H in the series are weaker cut downing agents than H.

A ( three ) A A The substances which are stronger oxidizing agents than H+ion are placed below H in the series.

( four ) A A The metals on the top ( holding high negative values of standard decrease potencies ) have the inclination to lose negatrons readily. These are active metals. The activity of metals lessenings from top to bottom. The non-metals on the underside ( holding high positive values of standard decrease potencies )

hold the inclination to accept negatrons readily. These are active non-metals. The activity of non-metals additions from top to bottom.

A Applications of Electrochemical series

A ( I ) A Reactivity of metals:

The activity of the metal depends on its inclination to lose negatron or negatrons, i.e. , inclination to organize cation ( M ” + ) . This inclination depends on the magnitude of standard decrease potency. The metal which has high negative value ( or smaller positive value ) of standard decrease possible readily loses the negatron or negatrons and is converted into cation. Such a metal is said to be chemically active.

The chemical responsiveness of metals lessenings from top to bottom in the series. The metal higher in the series is more active than the metal lower in the series. For illustration,

( a ) A Alkali metals and alkalic Earth metals holding high negative values of standard decrease potencies are chemically active. These react with cold H2O and germinate H. These readily dissolve in acids organizing matching salts and combine with those substances which accept negatrons.

( B ) A Metallic elements like Fe, Pb, Sn, Ni, Co, etc. , which lie a small down in the series do non respond with cold H2O but react with steam to germinate H.

( degree Celsius ) A Metallic elements like Cu, Ag and Au which lie below H are less reactive and do non germinate H from H2O.

A ( two ) A Electropositive character of metals:

The positively charged character besides depends on the inclination to lose negatron or negatrons. Like responsiveness, the positively charged character of metals lessenings from top to bottom in the electrochemical series. On the footing of standard decrease potency values, metals are divided into three groups:

( a ) A Strongly positively charged metals: Metallic elements holding standard decrease potency near about -2.0 V or more negative like alkali metals, alkalic Earth metals are strongly positively charged in nature.

( B ) A Reasonably positively charged metals: Metallic elements holding values of decrease potencies between 0.0 and about -2.0 V are reasonably positively charged. Al, Zn, Fe, Ni, Co, etc. , belong to this group.

( degree Celsius ) Weakly positively charged metals: The metals which are below H and possess positive values of decrease potencies are decrepit positively charged metals. Cu, Hg, Ag, etc. , belong to this group.

A ( three ) A A A A Displacement reactions:

( a ) A To foretell whether a given metal will displace another, from its salt solution. A metal higher in the series will displace the metal from its solution which is lower in the series, i.e. , the metal holding low standard decrease potenA­tial will displace the metal from its salt ‘s solution which has higher value of standard decrease potency. A metal higher in the series has greater inclination to supply negatrons to the cations of the metal to be precipitated.

( B ) A Displacement of one nonmetal from its salt solution by another nonmetal: A nonmetal higher in the series ( towards bottom side ) , i.e. , holding high value of decrease potency will displace another nonmetal with lower decrease possible i.e. , busying place above in the series. The nonmetal ‘s which possess high positive decrease potencies have the inclination to accept negatrons readily. These negatrons are provided by the ions of the nonmetal holding low value of decrease potency. Thus, Cl2A can displace Br and I from bromides and iodides.

Cl2A + 2KI — & gt ; 2KC1 + I2

21-A — & gt ; I2A + 2e-A A A A A A A A ( Oxidation )

Cl2A + 2e- — & gt ; 2C1-A A A A A ( Reduction )

[ The activity or negatively charged character or oxidising nature of the nonmetal additions as the value of decrease potency additions. ]

( degree Celsius ) A Displacement of H from dilute acids by metals: A The metal which can supply negatrons to H+A ions present in dilute acids for decrease, evolve H from dilute acids.

Mn — & gt ; A Mn ” +A + ne-A A A ( Oxidation )

2H+A + 2e-A — & gt ; H2A A A A ( Reduction )

The metal holding negative values of decrease potency possess the belongings of losing negatron or negatrons. Therefore, the metals busying top places in the electrochemical series readily liberate H from dilute acids and on falling in the series inclination to emancipate H gas from dilute acids lessenings.

The metals which are below H in electrochemical series like Cu, Hg, Au, Pt, etc. , do non germinate H from dilute acids.

( vitamin D ) Supplanting of H from H2O: A Iron and the metals above Fe are capable of emancipating H from H2O. The inclination decreases from top to bottom in electrochemical series. Alkali and alkalic Earth metals liberate H from cold H2O but Mg, Zn and Fe liberate H from hot H2O or steam.

A ( four ) Reducing power of metals:

Reducing nature depends on the inclination of losing negatron or negatrons. More the negative decrease potency, more is the inclination to lose negatron or negatrons. Therefore, cut downing nature decreases from top to bottom in the electrochemical series. The power of the cut downing agent additions as the standard decrease potency becomes more and more negative.

Sodium is a stronger cut downing agent than Zn and Zn is a stronger cut downing agent than Fe.

A ElementA A A A A A A A A A A A A A A A A A A A NaA A A A A A A A A A A A A A ZnA A A A A A A A A A A A A A A A A A Fe

Reduction potentialA A -2.71A A A A A A A A A A A -0.76A A A A A A A A A A A A A A A A A -0.44

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A — — — — — — — — — — — — — — — — — — — — — — & gt ;

Reducing nature lessenings

Alkali and alkalic Earth metals are strong cut downing agents.

A ( V ) Oxidising nature of nonmetals:

Oxidizing nature depends on the inclination to accept negatron or negatrons. More the value of decrease potency, higher is the inclination to accept negatron or electrons.A Therefore, oxidizing nature additions from top to bottom in the electrochemical series.A The strength of an oxidising agent additions as the value of decrease potency becomes more and more positive.

F2A ( Fluorine ) is a stronger oxidant than Cl2, Br2A and I2.

Cl2A ( Chlorine ) is a stronger oxidant than Br2A and I2.

Element A A A A A A A A A A A A A A A A A A A A A A A I2A A A A A A A A A A Br2A A A A A A A A A Cl2A A A A A A F2

Reduction potentialA A A A A +0.53A A +1.06A A A A +1.36A A +2.85

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A — — — — — — — — — — — — — — — — — — — – & gt ;

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A Oxidising nature additions

( six ) A Thermal stableness of metallic oxides:

The thermic stableness of the metal oxide depends on its positively charged nature. As the electropositivity decreases from top to bottom, the thermic stableness of the oxide besides decreases from top to bottom. The oxides of metals holding high positive decrease potencies are non stable towards heat. The metals which come below Cu signifier unstable oxides, i.e. , these are decomposed on warming.

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A Heat

Ag2OA A — — — — – & gt ; A 1/2A O2

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A 2 Ag

A A A A A A A A A A A A A A A A A A Heat

A A A A A A A A A A A A A A A A A A 2HgOA A — — — — — — & gt ; 1/2 O2

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A 2 Hg

A ( seven ) Products of electrolysis:

In instance two or more types of positive and negative ions are present in solution, during electrolysis certain ions are discharged or liberated at the electrodes in penchant to others.A In general, in such com request the ion which is stronger oxidizing agent ( high value of standard decrease potency ) is discharged foremost at the cathode.A The increasing order of deposition of few cations is:

K+ , Ca2+ , Na+ , Mg2+ , Al3+ , Zn2+ , Fe2+ , H+ , Cu2+ , Ag+ , Au3+

— — — — — — — — — — — — — — — — — — — — — — — — — — — — — — & gt ;

A A A A A A A A A A A A A Increasing order of deposition

Similarly, the anion which is stronger cut downing agent ( low value of standard decrease potency ) is liberated foremost at the anode.

The increasing order of discharge of few anions is:

A A A A A A A A A A A A A A A A A A A A A A A A A A A SO42- , NO3- , OH- , Cl- , Br- , I-

A A A A A A A A A A A A A A A A A A A A A A A — — — — — — — — — — — — — — — — — — – & gt ;

A A A A A A A A A A A A A A A A A A A A A A A A A Increasing order of discharge

Therefore, when an aqueous solution of NaCl incorporating Na+ , Cl- , H+A and OH ” ions is electrolysed, H+A ions are discharged at cathode and CF ions at the anode, i.e. , H2A is liberated at cathode and Cl at anode.

When an aqueous solution of CuS04A incorporating Cu2+ , , H+A and OH-A ions is electrolysed, Cu2+A ions are dis charged at cathode and OH-A ions at the anode.

A A Cu2+A + 2e-A — & gt ; CuA A A A A A A A A A A A A A A A A A A A A A A A A A A A A ( Cathodic reaction )

4OH-A — & gt ; O2A + 2H2O + 4e-A A A A A A A A A A A A A A A A A A A A A ( Anodic reaction )

Cu is deposited on cathode while 02A is liberated at anode.

( eight ) Corrosion of metals:

Corrosion is defined as the impairment of a substance because of its reaction with its environment. This is besides defined as the procedure by which metals have the inclination to travel back to their combined province, i.e. , contrary of extraction of metals.

Ordinary corrosion is a redox reaction by which metals are oxidised by O in presence of wet. Oxidation of metals occurs more readily at points of strain. Therefore, a steel nail foremost corrodes at the tip and caput. The terminal of a steel nail Acts of the Apostless as an anode where Fe is oxidised to Fe2+A ions.

A A A A A A A A A A A A A Fe — & gt ; Fe2A + 2e-A A A A A A ( Anode reaction )

The negatrons flow along the nail to countries incorporating im purenesss which act as cathodes where O is reduced to hydroxyl ions.

A A A A A A A A A A A A O2A + 2H2O + 4e-A — & gt ; 4OH-A ( Cathode reaction )

The overall reaction is

2Fe + OzA + 2H2O = & gt ; 2Fe ( OH ) 2

Fe ( OH ) 2A may be dehydrated to press oxide, FeO, or farther oxidised to Fe ( OH ) 3and so dehydrated to press rust, Fe203. Several methods for protection of metals against corrosion have been developed. The most widely used are ( one ) plating the metal with a thin bed of a less easy oxidised metal ( two ) leting a protective movie such as metal oxide ( three ) Galvani ing-steel is coated with Zn ( a more active metal ) .

A ( nine ) Extraction of metals:

A more positively charged metal can displace a less positively charged metal from its salt ‘s solution. This rule is applied for the extraction of Ag and Au by cyanide procedure. Silver from the solution incorporating Na argento nitrile, NaAg ( CN ) 2, can be obtained by the add-on of Zn as it is more electro-positive than Ag.

A A A A A A A A A A A A A A 2NaAg ( CN ) 2A + Zn — & gt ; Na2Zn ( CN ) 4A + 2Ag

standard redox possible

The values that we have merely quoted for the two cells are really the standard electrode potencies of the Mg2+ / Mg and Cu2+ / Cu systems.

The voltage measured when a metal / metal ion electrode is coupled to a H electrode under standard conditions is known as the standard electrode potency of that metal / metal ion combination. By convention, the H electrode is ever written as the left-hand electrode of the cell. That means that the mark of the electromotive force quoted ever gives you the mark of the metal electrode.

Standard electrode potency is given the symbol EA° .

Prediction For Occurrence of a Redox Chemical reaction

A Any redox reaction would happen spontaneously if the free energy alteration ( a?†G ) is negative. The free energy is related to cell emf in the undermentioned mode:

A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A A a?†Go = & gt ; – nFEo

where N is the figure of negatrons involved, F is the value of Faraday and Eo is the cel voltage. a?†G can be negative if Eo is positive.

A A A A A A A When Eo is positive, the cell reaction is self-generated and serves as a beginning of electrical energy.

To foretell whether a peculiar oxidation-reduction reaction will happen or non, compose down the redox reaction into two half reactions, one affecting decrease oxidization reaction and the other affecting decrease reaction. Write the oxidization reaction and decrease potency value for decrease reaction. Add these two values, if the algebraic summing up gives a positive value, the reaction will happen, otherwise non.

A

Arrhenius Theory of Electrolytic dissociation

In order to explicate the belongingss of electrolytic solutions, Arrhenius put away, in 1884, a comprehensive theory which is known as theory of electrolytic dissociation or ionic theory. The chief points of the theory are:

( I ) A A A A An electrolyte, when dissolved in H2O, breaks up into two types of charged atoms, one transporting a positive charge and the other a negative charge. These charged atoms are called ions. Positively charged ions are termed cations and negatively charged as anions.

AB — & gt ; A+A + B-

NaCl — & gt ; A Na+ + CL-

K2SO4 — & gt ; 2K++ SO42-

Electrolyte A A A A A A A A Ions

In its modern signifier, the theory assumes that solid electrolytes are composed of ions which are held together by electrostatic forces of attractive force. When an electrolyte is issolved in a dissolver, these forces are weakened and the electrolyte undergoes dissociation into ions. The ions are solvated.

A A+B- — & gt ; A+A + B-

or A A A A A A A A A A A A A+B-+ aq — & gt ; A A+ ( aq ) +B- ( aq )

A

( two ) A A A The procedure of splitting of the molecules into ions of an electrolyte is called ionisation. The fraction of the entire figure of molecules present in solution as ions is known as grade of ionisation or grade of dissociation. It is denoted by

I±= ( Number of molecules dissociated into ions ) / ( Entire figure of molecules )

It has been observed that all electrolytes do non ionise to the same extent. Some are about wholly ionized while others are feebly ionized. The grade of ionisation depends on a figure of factors ( see 12.6 ) .

A ( three ) A A A Ions nowadays in solution invariably re-unite to organize impersonal molecules and, therefore, there is a province of dynamic equilibrium between the ionized the ionised and non-ionised molecules, i.e. ,

A A A A A A A A A A A A A A A A A A A A A A ABA & lt ; — & gt ; A A+ + B-

Using the jurisprudence of mass action to above equilibrium

[ A+ ] [ B- ] / [ AB ] = & gt ; K

K is known as ionisation invariable. The electrolytes holding highA value of K are termed strong electrolytes and those holding low value of K as weak electrolytes.

A ( four ) A A When an electric current is passed through the electrolytic solution, the positive ions ( cations ) move towards cathode and the negative ions ( anions ) move towards anode and acquire discharged, i.e. , electrolysis occurs.

The ions are discharged ever in tantamount sums, no affair what their comparative velocities are.

ELECTROCHEMICAL CELL

A Electrochemical cell is a system or agreement in which two electrodes are fitted in the same electrolyte or in two different electrolytes which are joined by a salt span. Electrochemical cells are of two types:

( a ) A A A A A Electrolytic cell ( B ) A A A Galvanic or Gur cell

A Electrolytic cell

It is a device in which electrolysis ( chemical reaction affecting oxidization and decrease ) is carried out by utilizing electricity or in which transition of electrical energy into chemical energy is done.

A Galvanic or Gur cell

It is a device in which a oxidation-reduction reaction is used to change over chemical energy into electrical energy, i.e. , electricity can be obtained with the aid of oxidization and decrease reaction. The chemical reaction responsible for production of electricity takes topographic point in two separate compartments. Each compartment consists of a suited electrolyte solution and a metallic music director. The metallic music director Acts of the Apostless as an electrode. The compartments incorporating the electrode and the solution of the electrolyte are called half-cells. When the two compartments are connected by a salt span and electrodes are joined by a wire through galvanometer the electricity begins to flux. This is the simple signifier of Gur cell.

A EMF of A Galvanic Cell

Every galvanic or Gur cell is made up of two half-cells, the oxidization half-cell ( anode ) and the decrease half-cell ( cathode ) . The potencies of these half-cells are ever difA­ferent. On history of this difference in electrode potencies, the electric current moves from the electrode at higher potency to the electrode at lower potency, i.e. , from cathode to anode. The way of the flow of negatrons is from anode to cathode.

A A A A A A A A A A A A A A A A A A A A

Flow of electronsA

A A A A A A A A A Anode A & lt ; ============== & gt ; Cathode

A A A A A A A A A A A A A A A A A A A A A A A A A A Flow of current

The difference in potencies of the two half-cells is known as the electromotive force ( emf ) of the cell or cell potency.

The voltage of the cell or cell potency can be calculated from the values of electrode potencies of the two half-cells constitutA­ing the cell. The undermentioned three methods are in usage:

( I ) When oxidization potency of anode and decrease potenA­tial of cathode are taken into history:

ECello = & gt ; Oxidation potency of anode + Reduction potency of cathode

A A A A A A A A A A A A A A A A A A A A = & gt ; Eoxo ( anode ) + Eredo ( cathode )

( two ) When decrease potencies of both electrodes are taken into history:

A A A A A ECello = & gt ; Reduction potency of cathode – Reduction potency of anode

A A A A A A A A A A A A A A A A = & gt ; ECathodeo – EAnodeo A = & gt ; Erighto – Elefto

( three ) A A When oxidization potencies of both electrodes are taken into history:

A A A A A A = & gt ; Oxidation potency of anode – Oxidation potency of cathode

A A A ECelloA = & gt ; A Eoxo ( anode ) – Eredo ( cathode )

Daniell Cell

A It is designed to do usage of the self-generated oxidation-reduction reaction between Zn and cuprous ions to bring forth an electric current ( Fig.12.7 ) . It consists of two half-cells. The half-cells on the left contains a Zn metal electrode dipped in ZnSO4 solution.

A A A A A

The half-cell on the right consists of Cu metal electrode in a solution CuSO4. The half-cells are joined by a salt span that prevents the mechanical commixture of the solution.

A When the Zn and Cu electrodes are joined by wire, the undermentioned observations are made:

( I ) A There is a flow of electric current through the external circuit.

( two ) A The Zn rod loses its mass while the Cu rod additions in mass.

( three ) A The concentration of ZnSO4 solution additions while the concentration of Cu sulfate solution lessenings.

( four ) A The solutions in both the compartments remain electrically impersonal.

A During the transition if electric current through external circuit, negatrons flow from the Zn electrode to the Cu electrode. At the Zn electrode, the Zn metal is oxidized to zinc ions which go into the solution. The negatrons released at the electrode travel through the external circuit to the Cu electrode where they are used in the decrease of Cu2+ ions to metallic Cu which is deposited on the electrode. Therefore, the overall oxidation-reduction reaction is:

Zn ( s ) + Cu2+ A A Cu ( s ) + Zn2+ ( aq )

Therefore, indirect oxidation-reduction reaction leads to the production of electrical energy. At the Zn rod, oxidization occurs. It is the anode of the cell and is negatively charged piece at Cu electrode, decrease takes, topographic point ; it is the cathode of the cell and is positively charged.

A Thus, the above points van be summed up as:

( I ) A A A A Voltaic or Galvanic cell consists of two half-cells. The reactions happening in half-cells are called half-cell reactions. The half-cell in which oxidization taking topographic point in it is called oxidization half-cell and the reaction taking topographic point in it is called oxidization half-cell reaction. Similarly, the half-cell occurs is called decrease half-cell and the reaction taking topographic point in it is called decrease half-cell reaction.

( two ) A A A The electrode where oxidization occurs is called anode and the electrode where decrease occurs is termed cathode.

( three ) A A A Electrons flow from anode to cathode in the external circuit.

( four ) A A Chemical energy is converted into electrical energy.

( V ) A A A The net reaction is the amount of two half-cell reactions. The reaction is Daniel cell can be represented as.

The net reaction is the amount of two half-cell reactions. The reaction is Daniel cell can be represented as

A Oxidation half reaction, A A A A A A A A A A A Zn ( s ) — & gt ; A Zn2+ ( aq ) + 2e-

A Reduction half reaction, A A A A Cu2+ ( aq ) + 2e- — & gt ; A Cu ( s )

A A A — — — — — — — — — — — — — — — — — — — — — — — — — — — — — — -i?

A A A Net reaction A A A A A A A A A A A A A A A A Zn ( s ) + Cu2+ ( aq ) — & gt ; Zn2+ ( aq ) + Cu ( s )

VOLTAIC CELL

PRIMARY VOLTAIC CELL ( THE DRY CELL )

A In this cell, one time the chemicals have been consumed, farther reaction is non possible. It can non be regenerated by change by reversaling the current flow through the cell utilizing an external direct current beginning of electrical energy. The most common illustration of this type is dry cell.

The container of the dry cell is made of Zn which besides serves as one of the electrodes. The other electrode is a C rod in the Centre of the cell. The Zn container is lined with a porous paper. A damp mixture of ammonium chloride, manA­ganese dioxide, Zn chloride and a porous inert filler occupy the infinite between the paper lined zinc container and the C rod. The cell is sealed with a stuff like wax.

As the cell operates, the Zn is oxidised to Zn2+

A ZnA — – & gt ; A A Zn2+ + 2e-A A A A ( Anode reaction )

The negatrons are utilized at C rod ( cathode ) as the ammonium ions are reduced.

2NH4++2e- — & gt ; 2NH3 + H2 A A ( Cathode reaction )

The cell reaction is

Zn+ 2 NH4+ — – & gt ; A Zn2+ + 2NH3 + H2

Hydrogen is oxidized by MnO2 in the cell.

2MnO2 + H2 — – & gt ; 2MnO ( OH )

Ammonia produced at cathode combines with Zn ions to organize complex ion.

Zn2+ + 4NH3 — – & gt ; [ Zn ( NH3 ) 4 ] 2+

Ecell is 1.6 V

Alkaline prohibitionist cell is similar to ordinary dry cell. It contains potassium hydrated oxide. The reaction in alkalic dry cell are:

Zn + 2OH- — – & gt ; Zn ( OH ) 2 + 2e- A A A A A A A A A A A A A A A A A A ( Anode reaction )

A 2MnO2 + 2H2O + 2e- — – & gt ; 2MnO ( OH ) + 2OH- A A A A A A A A A A ( Cathode reaction )

Zn + 2MnO2 + 2H2O — – & gt ; Zn ( OH ) 2 + 2MnO ( OH ) A A A A A A ( Overall )

A Ecell is 1.5 V.

SECODARY VOLTAIC CELL ( LEAD STORAGE BATTERY )

A The cell in which original reactants are regenerated by go throughing direct current from external beginning, i.e. , it is re-charged, is called secondary cell. Lead storage battery is the illustration of this type.

It consists of a group of lead home bases bearing compressed squashy lead, jumping with a group of lead home bases bearing foliage dioxide, PbO2. These home bases are immersed in a solution of approximately 30 % H2SO4. When the cell discharge ; it operates as a Gur cell. The squashy lead is oxidized to Pb2+ ions and lead home bases get a negative charge.

A A A A A A A A A A A A A A A A A A A A A A A Pb — & gt ; Pb2+ + 2e- A A A A A A A A A A A A A A A A A A A A A ( Anode reaction )

Pb2+ ions combine with sulphate ions to organize indissoluble lead sulfate, PbSO4, which begins to surface lead electrode.

A A A A A A A A A A A A A A A A A A A A Pb2+ + SO42- — – & gt ; PbSO4 A A A A A A A A A A ( Precipitation )

A A A A A A A The negatrons are utilized at PbO2 electrode.

A A A A A A A A A A A A A A A PbO2 + 4H+ + 2e- — – & gt ; Pb2+ 2H2O A A A A A A A A A A A A A A ( Cathode reaction )

A A A A A A A A A A A A A A A A Pb2+ + SO42- — – & gt ; PbSO4 A A A A A A A A A A ( Precipitation )

A A A A A A A Overall cell reaction is:

A A A A A A A A A A A A A A A Pb + PbO2 + 4H+ + 2 SO42- — – & gt ; 2PbSO4 + 2H2O

A A Ecell is 2.041 V.

When a possible somewhat greater than the potency of battery is applied, the battery can be re-charged.

A A A A A A A A A A A A A A A A A A A 2PbSO4 + 2H2O — – & gt ; A Pb + PbO2 + 2H2SO4

After many repeated charge-discharge rhythms, some of the lead sulfate falls to the underside of the container, the sulfuric acid concentration remains low and the battery can non be recharged to the full.

A FUEL CELL

Fuel cells are another agencies by which chemical energy may be converted into electrical energy. The chief disadvantage of a primary cell is that it can present current for a short period merely. This is due to the fact that the measure of oxidizing agent and cut downing agent is limited. But the energy can be obtained indefinitely from a fuel cell every bit long as the outside supply of fuel is maintained. One of the illustrations is the hydrogen-oxygen fuel cell. The cell consists of three compartments separated by a porous electrode. Hydrogen gas is introduced into one compartment and O gas is fed into another compartment. These gases so diffuse easy through the electrodes and react with an electrolyte that is in the cardinal compartment. The electrodes are made of porous C and the electrolyte is a rosin incorporating concentrated aqueous Na hydrated oxide soluA­tion. Hydrogen is oxidised at anode and O is reduced at cathode. The overall cell reaction produces H2O. The reactions which occur are:

A AnodeA A [ H2 ( g ) + 2OH- ( aq ) A A A A A — – & gt ; A A A 2H2O ( cubic decimeter ) + 2e- ] A- 2

CathodeA A A A A A A A A A A O2 ( g ) + 2H2O ( cubic decimeter ) + 4e-A — – & gt ; A A 4OH- ( aq )

A A A A A A A A A A — — — — — — — — — — — — — — — — — — — — — — — — — — — — — — –

OverallA A 2H2 ( g ) + O2 ( g ) — – & gt ; A 2H20 ( cubic decimeter )

A This type of cells are used in space-crafts. Fuel cells are efficient and pollution free.

CONCENTRATIONS CELLS

If two home bases of the same metal are dipped individually into two solutions of the same electrolyte and are connected with a salt span, the whole agreement is found to move as a voltaic cell. In general, there are two types of concentration cells:

A ( I ) Electrode concentration cells:

In these cells, the possible difference is developed between two like electrodes at different concentrations dipped in the same solution of the electrolyte. For illustration, two H electrodes at different has force per unit area in the same solution of H ions constitute a cell of this type.

A A A A A A A A A A A A A ( Pt, H2 ( Pressure p1 ) ) /Anode |H+ | ( H2 ( Pressure p2 ) Pt ) /Cathode

If p1, p2 oxidization occurs at L.H.S. electrode and decrease occurs at R.H.S. electrode.

A A A A A A A A A A A Ecell = 0.0591/2 log ( p1/p2 ) A at 25o C

In the amalgam cells, two amalgams of the same metal at two different concentrations are interested in the same electrolyte solution.

A ( two ) Electrolyte concentration cells:

In these cells, electrodes are indistinguishable but these are immersed in solutions of the same electrolyte of different concentrations. The beginning of electrical energy in the cell is the inclination of the electrolyte to spread from a solution of higher concentration to that of lower concentration. With the termination of clip, the two concentrations tend to go equal. Therefore, at the start the voltage of the cell is maximal and it bit by bit falls to zero. Such a cell is repreA­sented in the undermentioned mode:

( C2 is greater than C1 ) .

A A A A A A A A A A A A A A A A A M|Mn+ ( C1 ) ||Mn+ ( C2 ) |M

or A A A A A A A A A A A A A ( Zn|Zn2+ ( C1 ) ) /Anode || ( Zn2+ ( C2 ) |Zn ) /Cathode

The voltage of the cell is given by the undermentioned look:

Ecell = 0.0591/n log C ( 2 ( R.H.S. ) ) /C ( 1 ( L.H.S. ) ) A at 25o C

The concentration cells are used to find the solubility of meagerly soluble salts, valency of the cation of the electrolyte and passage point of the two allotropic signifiers of a metal used as electrodes, etc.

Relation between Equilibrium invariable, Gibbs free energy and EMF of the cell

Concept of equilibrium in electrochemical cell

In an electrochemical cell a reversible oxidation-reduction procedure takes topographic point, e.g. , in Daniell cell:

A A A A A A A A A A A A A A A Zn ( s ) +Cu2+ ( aq ) & lt ; == & gt ; Zn2+ ( aq ) +Cu ( s )

( 1 ) A A At equilibrium mass action ratio becomes equal to equilibrium changeless,

A A A A A A A i.e. , A A A A A A A A A A Q = Ke

( 2 ) A A Oxidation potency of anode = -Reaction potency of cathode

A A A A emf = oxidization potency of anode + Reduction potency of cathode = 0

A A A A A A A Cell is to the full discharged

Harmonizing to Nernst equation:

A A A A A A A A A A A A A A A A A A A A A A A E = Eo – 0.0591/n log10 Q at 25o

At equilibrium, E = 0, Q = K

A A A A A A A A A A A A A A A A A A A A A A A A 0 = Eo 0.0591/n log10 K

A A A A A A A A A A A A A A A A A A A A A A A A K = Antilog [ ( nEo ) /0.0591 ]

A Work done by the cell

Let n faraday charge be taken out of a cell of emf E ; so work done by the cell will be calculated as:

Work = Charge x Potential A = nFE

Work done by the cell is equal to diminish in free energy.

-a?†G = nFE

Similarly, maximal gettable work from the cell will be

Wmax = nFEA°

where, Eo = standard voltage or criterion cell potency.

A -a?†G = nFE

A The relationship among K, a?†Go and Eo cell

A

Heat of reaction in an electrochemical cell

A Let N Faraday charge flows out of a cell of voltage E,

Then A A A A A A A A A -a?†G = nFEA A A A A A A A A A A A A A A A A A A A A … … ( I )

A Gibbs-Helmholtz equation from thermodynamics may be given as

a?†G = a?†H + T ( a?‚a?†G/a?‚T ) P A A A A A A A A A A A A A A A A A A … … ( two )

A From equation ( I ) and ( two ) we get

A -nFE = a?†H + TA ( a?‚ ( -nFE ) /a?‚T ) P = a?†H-nFT ( a?‚E/a?‚T ) P

A A a?†H = -nFE + nFT ( a?‚E/a?‚T ) P

Here ( a?‚E/a?‚T ) P = Temperature coefficient of cell

Case I: A When ( a?‚E/a?‚T ) P = 0, so a?†H=-nFE

Case II: When ( a?‚E/a?‚T ) & gt ; 0, so nFE & gt ; a?†H, i.e. , procedure inside the cell is endothermal.

Case III: A When ( a?‚E/a?‚T ) & lt ; 0, so nFE & lt ; a?†H, i.e. , procedure inside the cell is exothermal.

Nernst Equation

ELECTRODE AND CELL POTENTIALS /NERNST EQUATION

A A A A A A A The electrode potency and the voltage of the cell depend upon the nature of the electrode, temperature and the activities ( concentrations ) of the ions in solution. The fluctuation of electrode and cell potencies with concentration of ions in solution can be obtained from thermodynamic considerations. For a general reaction such as

A A A A A A A A A A A A A A A M1A + m2B… .. A n1X + n2Y + … .A A … … . ( I )

happening in the cell, the Gibbs free energy alteration is given by the equation

A A A A G = & gt ; a?†Go + 2.303RT log10 ( axn1 A- ayn2 ) / ( aAm1 A- aBm2 ) … … . ( two )

where ‘a ‘ represents the activities of reactants and merchandises under a given set of conditions and a?†Go refers to liberate energy alteration for the reaction when the assorted reactants and merchandises are present at standard conditions. The free energy alteration of a cell reaction is related to the electrical work that can be obtained from the cell, i.e. , a?†Go = -nFEcell and a?†Go = -nFEo. On replacing these values in Eq. ( two ) we get

-nFEcell = & gt ; -nFEo + 2.30eRT log10 ( axn1 A- ayn2 ) / ( aAm1 A- aBm2 ) A … … . ( three )

or A Ecell = & gt ; Ecello – 2.303RT/nF log10A ( axn1 A- ayn2 ) / ( aAm1 A- aBm2 ) … … . ( four )

This equation is known as Nearnst equation.

Puting the values of R= & gt ; 8.314 JK-1 mol-1, T = & gt ; 298 K and F= & gt ; 96500A C, Eq. ( four ) reduces to

E = & gt ; Eo – 0.0591/n log10 ( axn1 A- ayn2 ) / ( aAm1 A- aBm2 ) … … . ( V )

= & gt ; Eo A – 0.0591/n log10 ( [ Products ] ) / ( [ Reactants ] ) A A … … . ( six )

Potential of individual electrode ( Anode ) : A See the general oxidization reaction,

A A A A A A A M — & gt ; Mn+ + ne-

Using Nernst equation,

Eox = & gt ; Eoxo – 0.0591/nA log10 [ Mn+ ] / [ M ]

where Eox is the oxidization potency of the electrode ( anode ) , A is the standard oxidization potency of the electrode.

[ Note: The concentration of pure solids and liquids are taken as integrity. ]

A Eox = & gt ; Eoxo – 0.0591/nA log10 [ Mn+ ]

Let us see a Daniell cell to explicate the above equations. The concentrations of the electrolytes are non 1 M.

A A A A A A A Zn ( s ) +Cu2+ ( aq ) & lt ; = & gt ; A Zn2+ ( aq ) + Cu ( s )

A A A A A A A Zn ( s ) |Zn2+ ( aq ) ||Cu2+ ( aq ) |Cu

A Potential at zinc electrode ( Anode )

A A A A A A A Eox = & gt ; Eoxo – 0.0591/nA log10 [ Zn3+ ]

A Potential at Cu electrode ( Cathode )

A A A A A A A Ered = & gt ; Eredo – 0.0591/nA log10 [ Cu2+ ]

A Emf of the cell

A Ecell = & gt ; Eox + Ered = & gt ; ( Eoxo + Eredo ) – 0.0591/n [ Zn2+/Cu2+ ]

A The value of n = 2 for both zinc and Cu.

Let us see an illustration, in which the values of N for the two ions in the two half-cells are non same. For illustration, in the cell

A A A A A A A A A A A A A A A A A A A A A Cu|Cu2+||Ag+|Ag

A The cell reaction is

A A A A A A A A A A A A A A A A A A A A A A A Cu ( s ) + 2Ag+ — – & gt ; Cu2+ + 2Ag

The two half-cell reaction are:

A A A A A A A A A A A A A A A Cu — & gt ; A Cu2+ + 2e-

A A A A A A A A A A A A A A A Ag+ + e- — & gt ; Ag

The 2nd equation is multiplied by 2 to equilibrate the figure of negatrons.

A A A A A A A A A A A A A A A 2Ag+ + 2e- — & gt ; 2 Ag

Eox = & gt ; A Eoxo – 0.0591/2 log10 [ Cu2+ ]

Ered = & gt ; Eredo – 0.0591/2 log10 [ Ag+ ] 2

Ecell = Eox + Ered = & gt ; Eoxo – 0.0591/2 log10 [ Cu2+ ] / [ Ag+ ] 2A = & gt ;

Ecell = & gt ; 0.0591/2 log10 [ Cu2+ ] / [ Ag+ ] 2