electrochemistryclass12-151220102931

8/16/2019 electrochemistryclass12-151220102931

1/68


2/68

Electrochemistry

introduction

• Electrochemistry is the study ofproduction of electricity from energy

released during spontaneous chemicalreactions and the use of electricalenergy to bring about non-spontaneous

chemical transformations


3/68

Electrochemistry

Oxidation and Reduction

• What is reduced is the oxidizing agent.

! oxidizes "n by ta#ing electrons from it.• What is oxidized is the reducing agent."n reduces ! by gi$ing it electrons.


4/68

Electrochemistry

Electrochemical cells• %he cell that con$erts chemical energy to

electrical energy is called daniell&s cell. andhas an electrical potential e'ual to (.( )*hen concentration of "n+! and Cu+! ions isunity ,( mol dm 3. /uch a de$ice is called agal$anic or a $oltaic cell.

• 0f an external opposite potential is applied inthe gal$anic cell 12ig. 3.+,a and increased

slo*ly4 *e find that the reaction continues tota#e place till the opposing $oltage reachesthe $alue (.( ) *hen4 the reaction stopsaltogether and no current flo*s through thecell.

• 5ny further increase in the external potentialagain starts the reaction but in the opposite

direction . 0t no* functions as an electrolyticcell4 a de$ice for using electrical energy tocarry non-spontaneous chemical reactions.

"n,s ! Cu+!,a' 6 "n+!,a' ! Cu,s


5/68

Electrochemistry

2unctioning of 7aniell cell *henexternal $oltage Eext opposing

the cell potential is applied.


6/68

Electrochemistry

8al$anic cells• 0n this de$ice the 8ibbs energy of the spontaneous redox reaction is

con$erted into electrical *or# *hich may be used for running a motor orother electrical gadgets li#e heater4 fan4 geyser4 etc.

• 7aniell cell discussed earlier is one such cell in *hich the follo*ing redoxreaction occurs.

• "n,s ! Cu+! ,a' 6 "n+! ,a' ! Cu,s

• %his reaction is a combination of t*o half reactions *hose addition gi$es theo$erall cell reaction9

• ,i Cu+! ! +e 6 Cu,s ,reduction half reaction

• ,ii "n,s 6 "n+! ! +e ,oxidation half reaction

• %hese reactions occur in t*o different portions of the 7aniell cell. %hereduction half reaction occurs on the copper electrode *hile the oxidationhalf reaction occurs on the zinc electrode. %hese t*o portions of the cell arealso called half-cells or redox couples. %he copper electrode may be calledthe reduction half cell and the zinc electrode4 the oxidation half-cell.


7/68

Electrochemistry

• Each half- cell consists of a metallic electrode dipped into an electrolyte. %he t*o half-cells areconnected by a metallic *ire through a $oltmeter and a s*itch externally. %he electrolytes of thet*o half-cells are connected internally through a salt bridge.

• /ometimes4 both the electrodes dip in the same electrolyte solution and in such cases *e donot re'uire a salt bridge.

• 5t each electrode-electrolyte interface there is a tendency of metal ions from the solution todeposit on the metal electrode trying to ma#e it positi$ely charged. 5t the same time4 metalatoms of the electrode ha$e a tendency to go into the solution as ions and lea$e behind theelectrons at the electrode trying to ma#e it negati$ely charged.

• 5t e'uilibrium4 there is a separation of charges and depending on the tendencies of the t*oopposing reactions4 the electrode may be positi$ely or negati$ely charged *ith respect to thesolution.

• 5 potential difference de$elops bet*een the electrode and the electrolyte *hich is calledelectrode potential. When the concentrations of all the species in$ol$ed in a half-cell is unitythen the electrode potential is #no*n as standard electrode potential.

• 5ccording to 0:;5C con$ention4 standard reduction potentials are no* called standardelectrode potentials. 0n a gal$anic cell4 the half-cell in *hich oxidation ta#es place is calledanode and it has a negati$e potential *ith respect to the solution.

• %he other half-cell in *hich reduction ta#es place is called cathode and it has a positi$epotential *ith respect to the solution. %hus4 there exists a potential difference bet*een the t*oelectrodes and as soon as the s*itch is in the on position the electrons flo* from negati$eelectrode to positi$e electrode. %he direction of current flo* is opposite to that of electron flo*.


8/68

Electrochemistry

• %he potential difference bet*een the t*o electrodes of a gal$anic cell is called the cellpotential and is measured in $olts. %he cell potential is the difference bet*een the electrodepotentials ,reduction potentials of the cathode and anode.

• 0t is called the cell electromoti$e force ,emf of the cell *hen no current is dra*n throughthe cell. 0t is no* an accepted con$ention that *e #eep the anode on the left and the

cathode on the right *hile representing the gal$anic cell. 5 gal$anic cell is generallyrepresented by putting a $ertical line bet*een metal and electrolyte solution and putting adouble $ertical line bet*een the t*o electrolytes connected by a salt bridge. :nder thiscon$ention the emf of the cell is positi$e and is gi$en by the potential of the half- cell on theright hand side minus the potential of the half-cell on the left hand side i.e.4

E'uation9- Ecell < Eright Eleft%his is illustrated by the follo*ing cell reaction example 9

Cell reaction9Cu,s ! +5g!,a' 6 Cu+! ,a' ! + 5g,s

alf-cell reactions9 Cathode ,reduction9 +5g!,a' ! +e 6 +5g,s

5node ,oxidation9 Cu,s 6 Cu+!,a' ! +e

• 0t can be seen that the sum of the abo$e reactions leads to o$erall reaction in the cell andthat sil$er electrode acts as a cathode and copper electrode acts as an anode. %he cell canbe represented as9

• Cu,s=Cu+!,a'=5g!,a'=5g,s and *e ha$eEcell < Eright Eleft < E 5g! 5g ECu+! Cu


9/68

Electrochemistry

>easurement of electrode potential

• %he potential of indi$idual half-cell cannot be measured. We can measureonly the difference bet*een the t*o half-cell potentials that gi$es the emf ofthe cell. 0f *e arbitrarily choose the potential of one electrode ,half-cell thenthat of the other can be determined *ith respect to this.

• 5ccording to con$ention4 a half-cell called standard hydrogen electrode

represented by ;t,s +,g !

,a'4 is assigned a zero potential at alltemperatures corresponding to the reaction

! ,a'!e 6 (?++ .

• %he standard hydrogen electrode consists of a platinum electrode coated*ith platinum blac#. %he electrode is dipped in an acidic solution and pure

hydrogen gas is bubbled through it. %he concentration of both the reducedand oxidised forms of hydrogen is maintained at unity . %his implies that thepressure of hydrogen gas is one bar and the concentration of hydrogen ionin the solution is one molar.


10/68

Electrochemistry


11/68

Electrochemistry

• 0f the concentrations of the oxidised and the reduced forms ofthe species in the right hand half-cell are unity4 then the cellpotential is e'ual to standard electrode potential.

• /ometimes metals li#e platinum or gold are used as inert

electrodes. %hey do not participate in the reaction but pro$idetheir surface for oxidation or reduction reactions and for theconduction of electrons. 2or example4 ;t is used in the follo*inghalf-cells9ydrogen electrode9 ;t,s=+,g= !,a'

• With half-cell reaction9 ! ,a'! e 6 ( @ + +,g Aromineelectrode9 ;t,s=Ar +,a'= Ar ,a'

• With half-cell reaction9 ( @ + Ar +,a'! e 6 Ar ,a'


12/68

Electrochemistry

• 0f the standard electrode potential of an electrode is greater than zero thenits reduced form is more stable compared to hydrogen gas.

• /imilarly4 if the standard electrode potential is negati$e then hydrogen gas ismore stable than the reduced form of the species.

• 0t can be seen that the standard electrode potential for fluorine is the highestin the %able indicating that fluorine gas ,2+ has the maximum tendency toget reduced to fluoride ions ,2 and therefore fluorine gas is the strongestoxidising agent and fluoride ion is the *ea#est reducing agent.

• Bithium has the lo*est electrode potential indicating that lithium ion is the*ea#est oxidising agent *hile lithium metal is the most po*erful reducing

agent in an a'ueous solution. 0t is obser$ed that as *e go from top tobottom the standard electrode potential decreases and *ith this4 decreasesthe oxidising po*er of the species on the left and increases the reducingpo*er of the species on the right hand side of the reaction.

• Electrochemical cells are extensi$ely used for determining the p of

solutions4 solubility product4 e'uilibrium constant and other thermodynamicproperties and for potentiometric titrations.


13/68

Electrochemistry

/tandard Reduction ;otentials

Reductionpotentials for

manyelectrodesha$e been

measured and

tabulated.


14/68

Electrochemistry

Oxidizing and Reducing 5gents

• %he strongestoxidizers ha$e the

most positi$ereduction potentials.

• %he strongestreducers ha$e the

most negati$ereduction potentials.


15/68

Electrochemistry

Oxidizing and Reducing 5gents

%he greater thedifference bet*een

the t*o4 the greaterthe $oltage of thecell.


16/68

Electrochemistry

Electromoti$e 2orce ,emf

• %he potential difference bet*een theanode and cathode in a cell is called

the electromoti$e force ,emf.• 0t is also called the cell potential4 and is

designated E cell.


17/68

Electrochemistry

/tandard Cell ;otentials

%he cell potential at standard conditionscan be found through this e'uation9

E cell° < E red ,cathode E red ,anode° °

Aecause cell potential is based onthe potential energy per unit ofcharge4 it is an intensi$e property.


18/68

Electrochemistry

Cell ;otentials

• 2or the oxidation in this cell4

• 2or the reduction4

E red < D.F )°

E red < !D.3G )°


19/68

Electrochemistry

Cell ;otentials

E cell° < E red° ,cathode E red° ,anode

< !D.3G ) ,D.F )< !(.(D )


20/68

Electrochemistry

)oltaic Cells

0n spontaneousoxidation-reduction

,redox reactions4electrons aretransferred andenergy is released.


21/68

Electrochemistry

)oltaic Cells

• 5 typical cell loo#sli#e this.

• %he oxidation occursat the anode.

• %he reductionoccurs at thecathode.


22/68

Electrochemistry

)oltaic Cells

Once e$en oneelectron flo*s from

the anode to thecathode4 thecharges in eachbea#er *ould not be

balanced and theflo* of electrons*ould stop.


23/68

Electrochemistry

)oltaic Cells

• %herefore4 *e use asalt bridge4 usually a:-shaped tube that

contains a saltsolution4 to #eep thecharges balanced.Cations mo$e to*ard

the cathode. 5nions mo$e to*ard

the anode.


24/68

Electrochemistry

)oltaic Cells

• 0n the cell4 then4electrons lea$e theanode and flo*through the *ire to

the cathode.• 5s the electrons

lea$e the anode4 thecations formed

dissol$e into thesolution in theanode compartment.


25/68

Electrochemistry

)oltaic Cells

• 5s the electronsreach the cathode4cations in thecathode are

attracted to the no*negati$e cathode.

• %he electrons areta#en by the cation4

and the neutralmetal is depositedon the cathode.


26/68

Electrochemistry

Hernst&s e'uation

• We ha$e assumed in the pre$ioussection that the concentration of all the

species in$ol$ed in the electrodereaction is unity. %his need not beal*ays true. Hernst sho*ed that for the

electrode reaction9• >n!,a' ! ne 6 >,s


27/68

Electrochemistry

• 0t can be seen that E depends on the concentration of both


28/68

Electrochemistry

• 0t can be seen that E,cell. depends on the concentration of bothCu+! and "n+! ions. 0t increases *ith increase in theconcentration of Cu+! ions and decrease in the concentration of"n+! ions.

• Ay con$erting the natural logarithm to the base (D andsubstituting the $alues of R4 2 and % < +IJ K4 it reduces to

ED,cell. < E D,cell. D.DLI log 1"n+! + 1Cu+!

2or a general electrochemical HernstMs e'uation• a 5 ! b A 6 c C ! d 7 ,on passing it *ith ne-

ED,cell. < E D,cell. R% ln N n2

ED,cell. < E D,cell. R% ln 1Cc17d n2 15a1Ab


29/68

Electrochemistry

E'uilbirum constant from HernstMse'uation

• 0f the circuit in 7aniell celis closed then *e note that the reaction• "n,s ! Cu+!,a' 6 "n+!,a' ! Cu,s ta#es place and as time passes4 the

concentration of "n+! #eeps on increasing *hile the concentration of Cu+!

#eeps on decreasing.

• 5t the same time $oltage of the cell as read on the $oltmeter #eeps ondecreasing. 5fter some time4 *e shall note that there is no change in theconcentration of Cu+! and "n+! ions and at the same time4 $oltmeter gi$eszero reading. %his indicates that e'uilibrium has been attained. 0n thissituation the Hernst e'uation may be *ritten as9

ED,cell < D< E D,cell +.3D3 R% log 1"n+! +2 1Cu+!

%herefore4

E D,cell < +.3D3 R% log 1"n+! +2 1Cu+!

A t t ilib i


30/68

Electrochemistry

Aut at e'uilibrium4 1"n+! < Kc 1Cu+!

and at % < +IJK the abo$e e'uation can be*ritten as

ED,cell


31/68

Electrochemistry

• %hus4 the abo$e e'uation gi$es a relationshipbet*een e'uilibrium constant of the reactionand standard potential of the cell in *hich that

reaction ta#es place. %hus4 e'uilibriumconstants of the reaction4 difficult to measureother*ise4 can be calculated from the

corresponding E⊖ $alue of the cell.


32/68

Electrochemistry

Electrochemical cell and gibbsenergy of the reaction

• Electrical *or# done in one second is e'ual toelectrical potential multiplied by total chargepassed. 0f *e *ant to obtain maximum *or#

from a gal$anic cell then charge has to bepassed re$ersibly. %he re$ersible *or# done bya gal$anic cell is e'ual to decrease in its 8ibbs

energy and therefore4 if the emf of the cell is Eand n2 is the amount of charge passed andPr 8 is the 8ibbs energy of the reaction4 then

Pr 8 < n2E,cell

• We #no* that E,cell is an intensi$e parameter but Pr8 is an extensi$e


33/68

Electrochemistry

We #no* that E,cell is an intensi$e parameter but Pr 8 is an extensi$ethermodynamic property and the $alue depends on n.

%hus4 if *e *rite the reaction

"n,s ! Cu+! ,a' 6 "n+!,a' ! Cu,s

Pr 8 < +2E

,cell but *hen *e *rite the reaction as

+ "n ,s ! + Cu+! ,a' 6+ "n+!,a' ! +Cu,s

%hen the e'uation becomes

Pr8 < G2E,cell

• 0f the concentration of all the reacting species is unity4 then

Ecell < E ⊖cell and *e ha$e

Pr 8 < n2E⊖ ⊖cell

%hus4 from the measurement of E ⊖cell *e can obtain an important thermodynamic

'uantity4 Pr 8 4 standard 8ibbs energy of the reaction. 2rom the latter *e can⊖

calculate e'uilibrium constant by the e'uation9

Pr 8 < R% ln K.⊖


34/68

Electrochemistry

Conductance of electrolytic cells

• %he electrical resistance is represented by the symbol QR& and it is measuredin ohm , *hich in terms of /0 base units is e'ual to ,#g m+?,/3 5+.

• 0t can be measured *ith the help of a Wheatstone bridge. %he electricalresistance of any obSect is directly proportional to its length4 l4 and in$erselyproportional to its area of cross section4 5. %hat is4

*e also #no*4

%he constant of proportionality4 is called resisti$ity ,specific resistance.

• 0ts /0 units are ohm metre , m and 'uite often its submultiple4 ohmcentimetre , cm is also.

• ;hysically4 the resisti$ity for a substance is its resistance *hen it is onemetre long and its area of cross section is one m+. 0t can be seen that9

( m < (DD cm or ( cm < D.D( m

• %he in$erse of resistance R is called conductance 8


35/68

Electrochemistry

• %he in$erse of resistance4 R4 is called conductance4 84and *e ha$e the relation9 8 < (?R < 5?Tl


36/68

Electrochemistry

g y g palso depends on the temperature and pressure at *hich the measurements are made. >aterialsare classified into conductors4 insulators and semiconductors depending on the magnitude oftheir conducti$ity.

• >etals and their alloys ha$e $ery large conducti$ity and are #no*n as conductors. Certain non-metals li#e carbon-blac#4 graphite and some organic polymersV are also electronically

conducting.• /ubstances li#e glass4 ceramics4 etc.4 ha$ing $ery lo* conducti$ity are #no*n as insulators.

/ubstances li#e silicon4 doped silicon and gallium arsenide ha$ing conducti$ity bet*eenconductors and insulators are called semiconductors and are important electronic materials.Certain materials called superconductors by definition ha$e zero resisti$ity or infiniteconducti$ity.

• Earlier4 only metals and their alloys at $ery lo* temperatures ,D to (L K *ere #no*n to beha$eas superconductors4 but no*adays a number of ceramic materials and mixed oxides are also#no*n to sho* superconducti$ity at temperatures as high as (LD K.

• Electrical conductance through metals is called metallic or electronic conductance and is due tothe mo$ement of electrons. %he electronic conductance depends on

,i the nature and structure of the metal

,ii the number of $alence electrons per atom,iii temperature ,it decreases *ith increase of temperature.

• 5s the electrons enter at one end and go out through the other end4 the


37/68

Electrochemistry

5s the electrons enter at one end and go out through the other end4 thecomposition of the metallic conductor remains unchanged. %he mechanismof conductance through semiconductors is more complex.

• pure *ater has small amounts of hydrogen and hydroxyl ions ,(D >*hich lend it $ery lo* conducti$ity ,3.L (D L / m (. When electrolytes aredissol$ed in *ater4 they furnish their o*n ions in the solution hence itsconducti$ity also increases.

• %he conductance of electricity by ions present in the solutions is calledelectrolytic or ionic conductance. %he conducti$ity of electrolytic ,ionicsolutions depends on9

i the nature of the electrolyte addedii size of the ions produced and their sol$ation

iii the nature of the sol$ent and its $iscosityi$ concentration of the electrolyte

$ temperature ,it increases *ith the increase of temperature.• ;assage of direct current through ionic solution o$er a prolonged period can

lead to change in its composition due to electrochemical reactions

> t f d ti it f


38/68

Electrochemistry

>easurement of conducti$ity ofionic solutions

• 5 conducti$ity cell consists of t*o platinum electrodes coated*ith platinum blac# *hich has finely di$ided metallic ;t isdeposited on the electrodes electrochemically. %hese ha$e areaof cross section e'ual to Q5& and are separated by distance Ql&.

%herefore4 solution confined bet*een these electrodes is acolumn of length l and area of cross section 5. %he resistanceof such a column of solution is then gi$en by the e'uation9

R


39/68

Electrochemistry

%he 'uantity l?5 is called cell constant denoted by the symbol48V.

• 0t depends on the distance bet*een the electrodes and theirarea of cross-section and has the dimension of length ( and can

be calculated if *e #no* l and 5.• >easurement of l and 5 is not only incon$enient but also

unreliable. %he cell constant is usually determined bymeasuring the resistance of the cell containing a solution*hose conducti$ity is already #no*n.

• 2or this purpose4 *e generally use KCl solutions *hoseconducti$ity is #no*n accurately at $arious concentrations andat different temperatures. %he cell constant4 8V4 is then gi$en bythe e'uation9

8V< llA


40/68

Electrochemistry

of any solution.

• 0t consists of t*o resistances R3 and RG4 a $ariable resistance R( and the conducti$ity cell ha$ingthe un#no*n resistance R+.

• %he Wheatstone bridge is fed by an oscillator O *hich is a source of a.c. po*er in the audiofre'uency range LLD to LDDD cycles per second.

• ; is a suitable detector and the bridge is balanced *hen no current passes through thedetector. :nder these conditions9

:n#no*n resistance R+ < R( RG R3

• Once the cell constant and the resistance of the solution in the cell is determined4 the

conducti$ity of the solution is gi$en by the e'uation9U < cell constant< 8VR R

• %he conducti$ity of solutions of different electrolytes in the same sol$ent and at a gi$entemperature differs due to charge and size of the ions in *hich they dissociate4 theconcentration of ions or ease *ith *hich the ions mo$e under a potential gradient.

• 0t4 therefore4 becomes necessary to define a physically more meaningful 'uantity called molar

conducti$ity denoted by the symbol Xm. 0t is related to the conducti$ity of the solution by thee'uation9

>olar conducti$ity < Xm < U C


41/68

Electrochemistry

0n the abo$e e'uation4 if U is expressed in / m ( and the


42/68

Electrochemistry

0n the abo$e e'uation4 if U is expressed in / m and theconcentration4 c in mol m 3 then the units of Xm are in / m+ mol ( .

( / m+mol ( < (DG / cm+mol ( or ( / cm+mol ( < (D G / m+mol (.

) i ti f d ti it d l


43/68

Electrochemistry

)ariation of conducti$ity and molarconducti$ity *ith concentration

• Aoth conducti$ity and molar conducti$ity change *ith theconcentration of the electrolyte. Conducti$ity al*ays decreases*ith decrease in concentration both4 for *ea# and strongelectrolytes.

• %his can be explained by the fact that the number of ions perunit $olume that carry the current in a solution decreases ondilution.

• %he conducti$ity of a solution at any gi$en concentration is the

conductance of one unit $olume of solution #ept bet*een t*oplatinum electrodes *ith unit area of cross section and at adistance of unit length.

• 0f 5 and l are unity then 8Volar conducti$ity of a solution at a gi$en concentration is the


44/68

Electrochemistry

y gconductance of the $olume ) of solution containing one mole ofelectrolyte #ept bet*een t*o electrodes *ith area of crosssection 5 and distance of unit length. %herefore4

• Xm


45/68

Electrochemistry

$olume4 )4 of solution containing one mole of electrolyte also increases. 0t has beenfound that decrease in U on dilution of a solution is more than compensated byincrease in its $olume.

• 5t a gi$en concentration4 Xm can be defined as the conductance of the electrolyticsolution #ept bet*een the electrodes of a conducti$ity cell at unit distance but ha$ingarea of cross section large enough to accommodate sufficient $olume of solution thatcontains one mole of the electrolyte.

• When concentration approaches zero, the molar conductivity is known aslimiting molar conductivity and is represented by the symbol Emo.


46/68

Electrochemistry

/trong electrolytes

• 2or strong electrolytes4 Xm increases slo*ly *ith dilution andcan be represented by the e'uation9

Xm g/OG are #no*n as (-(4 +-( and +-+ electrolytesrespecti$ely. 5ll electrolytes of a particular type ha$e the same$alue for Q5&.

• Kohlrausch examined YmZ $alues for a number of strong


47/68

Electrochemistry

m gelectrolytes and obser$ed certain regularities.

• e noted that the difference in YmZ of the electrolytes Ha[ andK[ for any [ is nearly constant. 2or example at +IJ K9

• YZm ,KCl. YZm ,HaCl.< YZm ,KAr. YZm ,HaAr.< YZm ,K0. YZm ,Ha0. ≃ +3.G /cm+ mol (

and similarly it *as found thatYZm ,HaAr. YZm ,HaCl.< YZm ,KAr. YZm ,KCl. (.J / cm≃ + mol (

B f i d d t i ti f


48/68

Electrochemistry

Ba* of independent migration ofions

• %he la* states that limiting molar conducti$ityof an electrolyte can be represented as the sumof the indi$idual contributions of the anion and

cation of the electrolyte. %hus4 if \ZHa! and \ZCl are limiting molar conducti$ity of the sodiumand chloride ions respecti$ely4 then the limiting

molar conducti$ity for sodium chloride is gi$enby the e'uation9YZ,HaCl


49/68

Electrochemistry

g y g] anions then its limiting molar conducti$ity is gi$en by9

• YZ < ]!\D! ! ]-\D-

ere4 \D!

and \D

are the limiting molar conducti$ities of the cationand anion respecti$ely.


50/68

Electrochemistry

Wea# electrolytes• Wea# electrolytes li#e acetic acid ha$e lo*er degree of dissociation at

higher concentrations and hence for such electrolytes4 the change in Xm*ith dilution is due to increase in the degree of dissociation andconse'uently the number of ions in total $olume of solution that contains (mol of electrolyte.

• 0n such cases Yom increases steeply on dilution4 especially near lo*er

concentrations. %herefore4 YZm cannot be obtained by extrapolation of Xm tozero concentration.

• 5t infinite dilution i.e.4 concentration c 6 zero electrolyte dissociatescompletely ,^


51/68

Electrochemistry

Aut for *ea# electrolytes li#e acetic acid4 thee'uation of Ka becomes

5pplications of #ohlrausch&s la*

:sing Kohlrausch la* of independent migrationof ions4 it is possible to calculate Yom for anyelectrolyte from the \o of indi$idual ions.>oreo$er4 for *ea# electrolytes li#e acetic acid itis possible to determine the $alue of itsdissociation constant once *e #no* the YmZ andXm at a gi$en concentration c.


52/68

Electrochemistry

summary


53/68

Electrochemistry

Electrolytic cells and electrolysis• 0n an electrolytic cell external source of $oltage is used to bring about a chemical

reaction.• One of the simplest electrolytic cell consists of t*o copper strips dipping in an

a'ueous solution of copper sulphate.

• 0f a 7C $oltage is applied to the t*o electrodes4 then Cu+! ions discharge at thecathode ,negati$ely charged and the follo*ing reaction ta#es place9

• Cu+!,a' ! +e 6 Cu ,s Copper metal is deposited on the cathode. 5t the anode4copper is

• con$erted into Cu+! ions by the reaction9Cu,s 6 Cu+!,s ! +e

• %hus copper is dissol$ed ,oxidised at anode and deposited ,reduced at cathode.

%his is the basis for an industrial process in *hich impure copper is con$erted intocopper of high purity. %he impure copper is made an anode that dissol$es on passingcurrent and pure copper is deposited at the cathode.

• /odium and magnesium metals are produced by the electrolysis of their fusedchlorides and aluminium is produced by electrolysis of aluminium oxide in presenceof cryolite.

2 d & l f l t l i


54/68

Electrochemistry

2araday&s la* of electrolysis2araday&s t*o la*s of electrolysis are9•i) First Law: he amount o! chemical reaction which occursat any electrode during electrolysis by a current isproportional to the "uantity o! electricity passed through theelectrolyte #solution or melt).

•#ii) $econd Law: he amounts o! di!!erent substancesliberated by the same "uantity o! electricity passing throughthe electrolytic solution are proportional to their chemicale"uivalent weights #%tomic &ass o! &etal ' (umber o!

electrons re"uired to reduce the cation).%ccording to !araday *+t


55/68

Electrochemistry

;roducts of electrolysis• ;roducts of electrolysis depend on the nature of material being electrolysed

and the type of electrodes being used.• 0f the electrode is inert ,e.g9-platinum or gold4 it does not participate in the

chemical reaction and acts only as source or sin# for electrons. On the otherhand4 if the electrode is reacti$e4 it participates in the electrode reaction.%hus4 the products of electrolysis may be different for reacti$e and inert.

electrodes.• %he products of electrolysis depend on the different oxidising and reducing

species present in the electrolytic cell and their standard electrodepotentials.

• >oreo$er4 some of the electrochemical processes although feasible4 are so

slo* #inetically that at lo*er $oltages these do not seem to ta#e place andextra potential *hich is called o$erpotential has to be applied4 *hich ma#essuch process more difficult to occur.


56/68

Electrochemistry

Aatteries5ny battery ,actually it may ha$e one or morethan one cell connected in series or cell that *euse as a source of electrical energy is basically agal$anic cell *here the chemical energy of the

redox reaction is con$erted into electricalenergy .

%here are + types of batteries

a;rimary batteries

b/econdary batteries


57/68

Electrochemistry

Aatteries

;rimary batteries


58/68

Electrochemistry

;rimary batteries• 0n the primary batteries4 the reaction occurs only once and after

use o$er a period of time battery becomes dead and cannot be

reused again.%he most familiar example of this type is the dry cell is also #no*n

as Beclanche cell after its disco$erer *hich is used commonlyin our transistors and cloc#s.

• %he cell consists of a zinc container that also acts as anodeand the cathode is a carbon ,graphite rod surrounded bypo*dered manganese dioxide and carbon . %he space bet*eenthe electrodes is filled by a moist paste of ammonium chloride

,HGCl and zinc chloride ,"nCl+. %he electrode reactions arecomplex4 but they can be *ritten approximately as follo*s 9

• 5node9 "n,s 6 "n+! ! +e

• Cathode9 >nO+ ! HG! ! e- 6 >nO,O!H3

• 0n the reaction at cathode4 manganese is reduced from the ! G oxidationstate to the !3 state 5mmonia produced in the reaction forms a complex


59/68

Electrochemistry

state to the !3 state. 5mmonia produced in the reaction forms a complex*ith "n+! to gi$e 1"n ,H3G+!. %he cell has a potential of nearly (.L ).

• >ercury cell4 suitable for lo* current de$ices li#e hearing aids4 *atches4 etc.consists of zinc mercury amalgam as anode and a paste of gO andcarbon as the cathode. %he electrolyte is a paste of KO and "nO. %heelectrode reactions for the cell are gi$en belo*9

5node9

"n,g ! +O 6 "nO,s ! +O ! +e

Cathode9• gO ! +O ! +e 6 g,l ! +O

• %he o$erall reaction is represented by

• "n,g ! gO,s 6 "nO,s ! g,l

• %he cell potential is approximately (.3L ) and remains constant during itslife as the o$erall reaction does not in$ol$e any ion in solution *hoseconcentration can change during its life time.


60/68

Electrochemistry

5l#aline Aatteries


61/68

Electrochemistry

/econdary batteries• 5 secondary cell after use can be recharged by passing current through it in the

opposite direction so that it can be used again. 5 good secondary cell can undergo alarge number of discharging and charging cycles.

• %he most important secondary cell is the lead storage battery commonly used inautomobiles and in$ertors. 0t consists of a lead anode and a grid of lead pac#ed *ithlead dioxide ,;bO+ as cathode. 5 3J_ solution of sulphuric acid is used as anelectrolyte.

• %he cell reactions *hen the battery is in use are gi$en belo*9

5node9 ;b,s ! /OG+,a' 6 ;b/OG ,s ! +e

Cathode9 ;bO,s!/OG+-,a'!G!,a'!+e 6;b/OG ,s! ++O,l

• i.e.4 o$erall cell reaction consisting of cathode and anode reactions is9 ;b,s !;bO+,s ! ++/OG,a' 6 +;b/OG,s ! ++O,l

• On charging the battery the reaction is re$ersed and ;b/OG,s on anode and cathode

is con$erted into ;b and ;bO+4 respecti$ely.

5 lead storage battery ,secondary


62/68

Electrochemistry

5 lead storage battery ,secondarybattery

• 5nother important secondary cell is the nic#el-cadmium cell*hich has longer life than the lead storage cell but more


63/68

Electrochemistry

*hich has longer life than the lead storage cell but moreexpensi$e to manufacture.

• %he o$erall reaction during discharge is9

• Cd ,s ! +Hi,O3 ,s 6 CdO ,s ! +Hi,O+ ,s ! +O ,l

2uel cells


64/68

Electrochemistry

2uel cells• 8al$anic cells that are designed to con$ert the energy of combustion of fuels

li#e hydrogen4 methane4 methanol4 etc. directly into electrical energy are

called fuel cells.• One of the most successful fuel cells uses the reaction of hydrogen *ithoxygen to form *ater.

• 0n the cell4 hydrogen and oxygen are bubbled through porous carbonelectrodes into concentrated a'ueous sodium hydroxide solution. Catalysts

li#e finely di$ided platinum or palladium metal are incorporated into theelectrodes for increasing the rate of electrode reactions. %he electrodereactions are gi$en belo*9

• Cathode9 O+ ,g ! ++O,l ! Ge 6 GO ,a'

• 5node9 ++ ,g ! GO ,a' 6 G+O,l ! Ge

• O$erall reaction being9++,g ! O+,g 6 + +O,l

• %he cell runs continuously as long as the reactants are supplied. 2uel cellsproduce electricity *ith an efficiency of about D _ compared to thermalplants *hose efficiency is about GD_.


65/68

Electrochemistry

ydrogen 2uel Cells


66/68

Electrochemistry

corrosion• Corrosion slo*ly coats the surfaces of metallic obSects *ith oxides or other salts of

the metal.

• Corrosion of iron ,commonly #no*n as rusting occurs in presence of *ater and air.%he chemistry of corrosion is 'uite Complex but it may be considered essentially asan electrochemical phenomenon. 5t a particular spot of an obSect made of iron4oxidation ta#es place and that spot beha$es as anode and *e can *rite the reaction

• E D,2e+!?2e < D.GG )• 5node9 + 2e ,s 6 + 2e+! ! G e

• Electrons released at anodic spot mo$e through the metal and go to another spot onthe metal and reduce oxygen in presence of ! ,*hich is belie$ed to be a$ailablefrom +CO3 formed due to dissolution of carbon dioxide from air into *ater. ydrogen

ion in *ater may also be a$ailable due to dissolution of other acidic oxides from theatmosphere. %his spot beha$es as cathode *ith the reaction

Cathode9 O+,g!G!,a'!Ge- 6++O,l

• %he o$erall reaction being9+2e,s ! O+ ,g ! G!,a' 6 +2e +!,a' ! + +O ,l E ,+O


67/68

Electrochemistry

*hich come out as rust in the form of hydrated ferric oxide ,2e+O3. x+Oand *ith further production of hydrogen ions.

• ;re$ention of corrosion is of prime importance. 0t not only sa$es money but

also helps in pre$enting accidents such as a bridge collapse or failure of a#ey component due to corrosion.

• One of the simplest methods of pre$enting corrosion is to pre$ent thesurface of the metallic obSect to come in contact *ith atmosphere. %his canbe done by co$ering the surface *ith paint or by some chemicals ,e.g.bisphenol.

• 5nother simple method is to co$er the surface by other metals ,/n4 "n4 etc.that are inert or react to sa$e the obSect. 5n electrochemical method is topro$ide a sacrificial electrode of another metal ,li#e >g4 "n4 etc. *hichcorrodes itself but sa$es the obSect.


68/68

Electrochemistry

o* to ;re$ent Corrosion `

Documents

electrochemistryclass12-151220102931