26772095 matriculation chemistry electrochemistry

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Chapter 10

ElectrochemistElectrochemist

ryry



Electrochemistry

Is the study of the relationship between electricity

and chemical reaction

Chemical reactions involved in electrochemistry are :Chemical reactions involved in electrochemistry are :

Oxidation

ReductionREDOX REACTION

One type of reaction cannot occur withoutOne type of reaction cannot occur without

the other.the other.



REDUCTION

gain of electron

Oxidation no. decrease

Reaction at cathode

Remember…RED CATRED CAT

= REDREDuction at CAT CAT hode

Example:

Cu2+ + 2e- Cu

Oxidation no. ↓

OXIDATION

loss of electron

Oxidation no. increase

Reaction at anode

Example:

Mg Mg2+ + 2e-

Oxidation no.

REDOX Reaction



Reduction :

Oxidation :

Cu2+(aq) + 2e- → Cu(s)

n(s) → n2+(aq) + 2e- Half-cellreaction

Cu2+(aq) + n(s) → Cu(s) + n2+

(aq)

Example

Overall cell

reaction :

Electrochemical reaction consists of reduction

and oxidation.These two reactions are called ‘half-cell

reactions’

The combination of 2 half reactions are called

‘cell reaction’



CellsThere are 2 type of cells

ElectrochemicalCells

ElectrolyticCells

where chemical reaction

produces electricity

Uses electricity to

produce chemical

reaction

ChemicalEnerg

ElectricalEnerg

ElectricalEnerg

ChemicalEnerg

Also called;

alvanic cell or !oltaic cell



Component and Operationof al!anic cell

Consists of :

"# $n metal in an a%ueous solution of $n2&

2# Cu metal in an a%ueous solution of Cu2&

- !he 2 metals are connected "# a $ire- !he 2 containers are connected "# a salt "ridge.- % &oltmeter is used to detect &oltage generated.



nelectrode

Cuelectrode

'alt

"ridge

Zn2+ Cu2+

oltmeter

Galvanic cell

n'O(aq)solution

Cu'O(aq)solution



"hat happens at the #incelectrode $

inc is more electropositi&e than copper.

!endenc# to release electrons: n * Cu.

n (s) → n2+ (aq) + 2e-

inc dissol&es. Oxidation occurs at the n electrode. n2+ ions enter n'O

) solution.

n is the &e electrode since it is a source of

electrons anode.



Cu2+ (aq) + 2e- → Cu (s)

Copper is deposited. Reduction occurs at the Cu electrode.

Cu is the +&e electrode cathode

"hat happens at the copperelectrode $

!he electron from the n metal mo&es out through the $ire enter the Cu metal

Cu2+ ions from the solution accept electrons.



Reactions ,n&ol&ed:

%node :Cathode :

n (s) → n2+

(aq) + 2e-

Cu2+ (aq) + 2e- → Cu (s)

n (s) + Cu2+ (aq)→ n2+ (aq) + Cu (s)O&erall cellreaction :



%&nctions

'alt "ridge helps to maintain electrical neutralit#

Completes the circuit "# allo$ing ions carr#ing chargeto mo&e from one half-cell to the other.

'alt brid(e

%n in&erted tu"e containing a gel%n in&erted tu"e containing a gelpermeated $ith solution of anpermeated $ith solution of an inert electrol#te such as Cl/ 0aelectrol#te such as Cl/ 0a22'O'O/ 01/ 010O0O..



'hat happened if there is no salt brid(e)

Zn Cu

eZn2+ Cu2+

V

e

e

e

ee

ZnSO4(aq !uSO4(aq

%s the 3inc rod dissol&es/ the concentration of n2+

in the left "ea4er increase. !he reaction stops "ecause the nett increase in

positi&e charge is not neutrali3ed.This excess char"e build#u$ can be reduced b% addin" a salt

brid"e



Salt brid"e

(&!l

Zn2+

- +Zn

ZnSO4(aq !uSO4(aq

Cu

ee

n (s) → n2+ (aq) + 2e- Cu2+ (aq) + 2e- → Cu (s)

%0O5E (-) C%!1O5E (+)

Cu2+

E = +!.!" #



*ow does the cell maintains its electrical neutrality)

$eft Cell Right Cell

Cl- ions from salt"ridge mo&e into n

half cell

+ ions from salt"ridge mo&e into Cu

half cell

Electrical neutralit# is maintained

n (s) → n2+ (aq) + 2e- Cu2+ (aq) + 2e- → Cu (s)

n2+ ions enter the solution.Causing an o&erall excess of

t&e charge.

Cu 2+ ions lea&e the solution.Causing an o&erall excess of

-&e charge.



Electrochemical !ells

'.2

s$ontaneous

redox reaction

anodeoxidation

cathodereduction

half



Zn (s ) !u2) (aq !u (s ) Zn2) (aq

Zn (s * Zn2) (aq ** !u2) (aq * !u (s

anode cathode

lso can !e represented as:

Cell notation

6hase "oundar#

'alt "ridge

Cell notation



Exercise

7or the cell "elo$/ $rite the reaction at anodeand cathode and also the o&erall cell reaction.

n (s) 8 n2+ (aq) 88 Cr+ (aq) 8 Cr (s)

n(s) n→ 2+ (aq) + 2e-

Cr+ (aq) + e- Cr(s)→

n(s) + 2Cr+ (aq) + → n2+ (aq) +2Cr(s)

Cathode :

%node :

O&erall cellreaction:

9 2

9

Cell notation

2 2e

e-



The difference in electrical $otential between the anode

and cathode is called:

+ cell voltage

+ electromotive force (emf)

+ cell potential

cts as ‘electrical pressure’ that

pushes electron throu"h the #ire$

measured "# a &oltmeter



Electrode %otential

measure of the a!ilit& of a half-cell to attract electrons to#ards it$

Cu2+ 'a() + 2e * Cu's) Eored +,$. /

Zn2+ 'a() + 2e * Zn's) Eored -,$01 /

tandard reduction potential of copper half-cell is

more positive compared to 3inc$

tandard reduction

potential+4he more positive the half-cell’s electrode potential5 the

stron"er the attraction for electrons$

Tendenc% for reduction ,

(cathode

Zinc half-cell !ecomes anode$



!ell -otential (Eocell Eo

catode / Eoanode

)0.14 (#0.3

)'.' 5

!u2) (aq ) 2e 6 !u(s Eored )0.14 5

Zn2) (aq ) 2e 6 Zn(s Eored #0.3 5



n2+ (aq) + 2e- → n (s)

Cu2+ (aq) + 2e- → Cu (s)

E" = -".%&#

E" = +".'(#

E"cell = E"

cathode - E"anode

= +".'( ) *-".%&

= +!.!" #or E"cell = E"

red + E"ox

= +".'( + *+".%&

= +!.!" #

Change the sign

1alf-cell equation at:

%node :

Cathode :

n (s) → n2+ (aq) + 2e-

Cu2+

(aq) + 2e-

→

Cu (s)



tandard Electrode %otentials 'Eo)

at 26oC5 the pressure is 7 atm 'for

"ases)5 and the concentration of

electrol&te is 78$

measure of the a!ilit& of half-cell

to attract electrons to#ards it

4he si"n of E, chan"es #hen the

reaction is reversed

+Chan"in" the stoichiometric

coefficients of a half-cell reaction

does not chan"e the value of E,



7or example:

Cl2*g + 2e- 2Cl-*a, E" = +!.'& #

Cl2*g + e- Cl-*a, E" = +!.'& #

Cl-*a, Cl2*g + e- E" = -!.'& #



tandard 9&dro"en Electrode '9E)

8ade up of a platinum electrode5 immersed in an

a(ueous solution of 9+ '7 8) and !u!!led #ithh&dro"en "as at 7 atm pressure5 and temperature at

26oC

telectrodeH+ *a,

! M

H2 ga/at ! atm

4he standard reduction of 9E is , /



Standard Electrode -otentials

Zn (s * Zn2)

(' M ** 7)

(' M * 72 (' atm * -t (s

2e# ) 27) (aq8' M 72 ("8' atm

Zn (s Zn2) (' M ) 2e# 9node (oxidation:

!athode (reduction:

Zn (s ) 27) (aq8' M Zn2) (aq ) 72 ("8' atmCell reaction

0.3 5 0 - E Zn /Zn 0

2)

E Zn /Zn #0.3 50 2)

Zn2)

) 2e#

Zn E 0

#0.3 5

E 0 E H/H

- E Zn /Zn cell

0 0 ) 2)

2



'tandard reduction potential of Copper half cell ismeasured "# setting up the electrochemical

cell as "elo$.

92 '") 26o

C7 atm$

Cu

Cu2

CuO.'a()

78

9+

'a()78

/

-t

E, , + -

+

-



-t (s * 72 (' atm * 7) (' M ** !u2) (' M * !u (s

2e# ) !u2) (' M !u (s

72 (' atm 27) ) 2e# 9node (oxidation:

!athode (reduction:

72 ) !u2) !u (s ) 27)

E 0 E cathode - E anodecell 0 0

E cell E Cu /Cu E H /H 2) )

2

0 0 0

0.14 E Cu /Cu # 00 2)

E Cu /Cu 0.14 52)0



!he direction of half-reaction of '1E depends on theother

half-cell connected on it.!he cell notation for '1E is either:

t*/ 0 H2*g 0 H+ *a, 1hen it i/ anode

H+*a, 0 H2*g 0 t*/ 1hen it i/ cathode

,n either case/ E; of '1E remains ;



n2+ (aq) + 2e- → n (s)

Cu2+ (aq) + 2e- → Cu (s)

E" = -".%&#

E" = +".'(#

E"cell = E"

cathode - E"anode

= +".'( ) *-".%&

= +!.!" #or E"cell = E"

red + E"ox

= +".'( + *+".%&

= +!.!" #

Change the sign

1alf-cell equation at:

%node :

Cathode :

n (s) → n2+ (aq) + 2e-

Cu2+

(aq) + 2e-

→

Cu (s)



%t standard-state condition

E 0cell = E 0red + E 0ox

E 0cell = E 0cathode - E 0anode

or

E i



ExerciseCalculate the standard cell potential of the follo$ing electrochemical cell.

Co(s) 8 Co2+ (aq) 88 %g+(aq) 8 %g(s)

AnswerCathode (Red) :

%node (Ox) : Co(s) → Co2+ (aq) + 2e-

%g+(aq) + e- → %g(aq) E; < +;.=;

E;ox < +;.2=

E"cell = E"

cathode - E"anode

= +"." ) *-".2

= +!." #

%g+(aq) + e- → %g(aq)

Co2+

(aq) + 2e-→ !o(s E

;

< -;.2=

E; < +;.=;

Refer to the list of 'tandard Reduction 6otential:



Oxidation agent left of the half cell e,uation)

Reduction agent right of the half cell e,uation)

Example :

Oxidationagent

Reducingagent

Refer to the list of 'tandard Reduction 6otential:

%g+ (aq) + e- %g (s)→ E; < +;.=;

Cu2+ (aq) + 2e- Cu (s)→ E; < +;. 0i2+ (aq) + 2e- 0i (s)→ E; < -;.2> ,ncrease

strength asreducing

agent

The more +ve the value of E0 the stronger→ the oxidizing agent

The more -ve the value of E0 the stronger→ the reducing agent



%rrange the elements in order of increasingstrength of reducing agents

3'+ + 'e- 3 E" = -!.&& #

42+

+ 2e-

4 E"

= -2.% #

$2+ + 2e- $ E" = +".5 #

Answer :$ 6 3 6 4

Exercise

Exam$le



!alculate the E0 cell for the reaction ::

"($ * "2)(a; ** Sn4)

(a; 8Sn2)(a; * -t($

<iven :

"2)(a; ) 2e6 "($ E= #2.1> 5

Sn4)(a; ) 2e 6 Sn2)

(a; E= )0.'? 5

Oxidation : "($ 6 "2)(a; ) 2e Eo

ox )2.1> 5

@eduction : Sn4)(a; ) 2e 6 Sn2)

(a; Eo )0.'? 5

"($ ) Sn

4)

(a; 6 "

2)

(a; ) Sn

2)

(a; Ecell )2.?1 5

Exam$le

E cell E o red + E o ox

E 0 cell E cathode - E anode

+,$76- '-2$;)

+2$6/

)2.1> ) 0.'?

)2.?1 5



Exercis

e% cell is set up "et$een a chlorine electrode and ah#drogen electrode

(a) 5ra$ a diagram to sho$ the apparatus and chemicals

used.(a) 5iscuss the chemical reactions occurring in theelectrochemical cell.

6t 8 12(g/ ? atm) 8 1+(aq/ ?@) 88 Cl2(g/ ?atm) 8 Cl-(aq/ ?@) 8 6t

E"cell

= +!.'& #



Answer

92 '")5

7 atm$

%t

9+'a()5 78

%t

E,cell 7$1/

- +

/- + Cl2 '")5

7 atm$

Cl-'a()5 78



7$ ho# the process occur at anode and

cathode

2$ Overall reaction

- 9alf-cell reaction

Answer



Answer Reduction (cathode)

Cl2 *g + 2e- 2Cl- *a,

Oxidation(anode)

H2 *g 2H+

*a, + 2e-

E

o

cell <+?.

E, ,

Eocell < Eo

cathode - E;anode

+?. < Eocathode ;

E;cathode < +?. 'o the standard reduction potential forCl2 is:

Eo < +?.



+ S$ontaneous A Bon#S$ontaneous

reactions # @edox reaction is s$ontaneous when

Ecell is )ve.

# Bon s$ontaneous is when Ecell is ve.

E= cell 0 The reaction is at equilibrium

%redict #hether the follo#in" reactions occur



n.+ <n2+

%redict #hether the follo#in" reactions occur

spontaneousl& or non-spontaneousl& under standard

condition$

Zn + n.+ * n2+ + Zn2+

4he t#o half-cells involved are:-

nod : Zn * Zn2+ + 2e Eo

ox +,$01 /

Cathode: n.+ + 2e * n2+ Eo +,$76 /

Zn + n.+ ** Zn2+ + n2+

Eocell Eo

+,$=7 /

spontaneous

Zn<Zn2+* Eo

+,$76 > '-,$01 )

Eocell Eo

red + Eoox

'+,$76) + ',$01)

+,$=7 /Or

n.+ <n2+Eo +,$76/

Eo

Zn<Zn2+ - ,$01/$



-b2)(aq ) 2!l#(aq 6 -b(s ) !l2("

-redict : S$ontaneous or non#

s$ontaneousC

Reduction



%!2+ 'a() + 2Cl-'a() * %!'s) + Cl2'")%!2+ 'a() + 2Cl-'a() * %!'s) + Cl2'")

cathode: %!2+ 'a() + 2e * %!'s)

anode: 2Cl-'a() * Cl2'") + 2e

%!2+ 'a() + 2Cl-'a() * %!'s) + Cl2' ")

Eo -,$7 /

Eoox -7$1

Eocell Eo

red + Eo

ox

Oxidation

'-7$1) + '-,$7)

-7$.; /

?on-spontaneous

?o Reaction

Example :



p

29"(s 29")(aq ) 2e E=

ox # 0.>0 5

Dr 2(aq ) 2e 2Dr #(aq E=

)'.03 5

29"(s ) Dr 2(aq 29")(aq ) 2Dr #(aq Esel

) 0.23 5

The reaction is spontaneous

Answer :

Predict whether the following reactions occurspontaneously :

29"(s ) Dr 2(aq 29")(aq ) 2Dr #(aq

E Ag /Ag )0.> 5)0

2E Br /Br )'.03 50

-

standard reduction

$otential



Exercise

% cell consists of sil&er and tin in a solution of ? @sil&er ions and tin (,,) ions. 5etermine the spontaneit#of the reaction and calculate the cell &oltage of thisreaction.

%g+ (aq) + e- %g (s)→

'n2+ (aq) + 2e- 'n (s)→

E; < +;.=; E; < -;.?

(cathode)

(anode)

E;

cell < E;

cathode - E;

anode

< +;.=; (-;.?)< +;.A

E;

cell < +&e ( reaction is spontaneous)

?ernst e(uation



?ernst e(uation

ernst e!uation can "e used to calculate the E cell

for any chosen concentration :

Ecell Eocell @T ln $roduct Fx

nG reactantF%

9t 2> & and @ >.1'4 H &#' mol#' 8 ' G ?00 !

Ecell Eocell 0.02?3 $roduct Fx

n reactantF%

2.101 lo"



Ecell Eocell 0.0?2 $roduct Fx

n reactantF%

lo"

Ecell Eocell 0.0?2

n

lo" Q

n # no of e- that are involvedQ # reaction !uotient

$ product %x

$ reactant%yQ



E*ample1

&alculate the Ecell for the following cell

Zn(s I Zn2) (aq8 0.02 II !u2) (aq8 0.40 I !u(s

Answer

'n(s) + &u*+ (a!) 'n*+ (a!) + &u(s)

Eocell Eo

red ) Eoox J

)0.14 5 ) 0.3 5

)'.'0 5

Eocell Eo

cathode Eoanode

)0.14 5 # (# 0.3 5

)'.'0 5



E # Eo 0,0.* log $ 'n*+ %

n $ &u*+ %

E # +/,/0 0,0.* log (0,0* )

* ( 0,10)

# +/,/0 (-0,023)

# +/,/2.

At e!uili"rium:



At e!uili"rium:

4 o net reaction occur (Q#K )

4 Ecell # 0

Ecell Eocell 0.0?2

n

lo" K

0 Eocell 0.0?2

n

lo" K

E=cell 0.0?2

nlo" K



E*ample+&alculate the e!uili"rium constant (K ) for the

following reaction,

&u(s) + *Ag+(a5) &u*+

(a5) + *Ag(s)

At e!uili"rium6 E cell # 0

E ocell # E o cathode - E o anode

# +0,30 ( +0,21)

# +0,17

Answer



Ecell # Eocell 0,0.* log K

*

0 # 0,17 0,0.* log K

*

0,0.* log K # 0,17

*

log K # /,1

K # 2,178 x /0/

El l i



ElectrolysisElectrol#sis is a chemical process that uses electricit#

for a non-spontaneous redox reaction to occur.'uch reactions ta4e place in electroltic cell/.

Electrolytic Cell ,t is made up of 2 electrodes immersed in an

electrol#te.

% direct current is passed through the electrol#tefrom an external source.

@olten salt and aqueous ionic solution are commonl#

used as electrol#tes.



Anion Cation

Oxidation Reduction

Electrolytic Cell

Electrolyte

(M+

X-

)X-,OH- M+,H+

+ -



An

ode

Cat

hod

e

6ositi&e electrode

!he electrode $hich is connected to the

positi&e terminal of the "atter# Oxidation ta4es place

0egati&e electrode

!he electrode $hich is connected to thenegati&e terminal of the "atter#

Reduction ta4es place

Electron/ flo1 from anode to cathode



Electrode

as circuit connectors

as sites for the precipitation of insolu"leproducts

example: 6latinum / Braphite (inert electrode)

Electrolyte

a liquid that conducts electricit# due to the presence of +&e and &e ions must "e in molten state or in aqueous

solution so that the ions can mo&e freel#example: Cl(l)/ 1Cl(aq)/ C1COO1(aq)

parison bet,een an electrochemical



parison bet,een an electrochemicaland an electrolytic cell

Cathode%node

- +e- e-

%node Cathode

+ -

e- e-

+ -

Electrolytic Cell Electrochemical Cell

El t l ti C ll El t h i l C ll



Electrolytic Cell Electrochemical Cell

Cathode < negati&e

%node < positi&e

Cathode < positi&e

%node < negati&e

0on-spontaneous redoxreaction requires energ#

to dri&e it

'pontaneous redoxreaction releases energ#

Oxidation occurs at anode/ reduction occursat cathode

%nions mo&e to$ards anode/ cations mo&eto$ards cathode. Electrons flo$ from anode to cathode in anexternal circuit.

7imilaritie/8



Electrolysis of molten salt

Electrol#sis of molten salt requires high temp. Electrol#sis of molten 0aCl

Cation : 0a+ %nion : Cl-

Anode 8

Cathode 8 +a& ,l# & e- +a ,s#→

Cl- ,l# Cl→ 2,(# & 2e-

O9erall 8 2+a& ,l# & 2Cl-,l# Cl→ 2,(# & 2+a,s#

2+a& ,l# & 2e- 2+a ,s#→

El t l sis f m lt 0 Cl i s s di m m t l



Electrol#sis of molten 0aCl gi&es sodium metaldeposited at cathode and chlorine gas e&ol&ed atanode.

Electrol#sis of molten 0aCl is industriall# important.!he industrial cell is called 5o$ns CellD

El t l i f A ' lt



Electrolysis of A-&eo&s 'alt Electrol#sis of aqueous salt is more complex

"ecause the presence of $ater. %queous salt solutions contains anion/ cation and $ater. ater is an electro-acti&e su"stance that ma# "e

oxidised or reduced in the process depending on

the condition of electrol#sis.

Reduction :

Oxidation :

2*2O ,l# & 2e- *2 ,(# & 2O*- ,a%#

2*2O ,l# *& ,a%# & O2 ,(# & e-

E; < -;.=

E; < -?.2



%redictin" the products of electrol&sis

Factors influencing the products :

7$ Reduction<oxidation potential of the

species in electrol&te2$ Concentrations of ions

$ 4&pes of electrodes used > active or

inert

Electrolysis of A-&eo&s NaCl



Electrolysis of A-&eo&s NaCl

!he electrol#sis of aqueous 0aCl depends on theconcentration of electrol#te.

0aCl aqueous solution contains 0a+ cation/ Cl- anion

and $ater molecules On electrol#sis/

the cathode attracts 0a+ ion and 12O molecules

the anode attracts Cl-

ion and 12O molecules

Electrolysis of dil&ted NaCl sol&tion



Cathode

2*2O ,l# & 2e- *2 ,(# & 2O*- ,a%# E; < -;.= +a

&

,a%# & e-

+a ,s# E; < -2.F?

E; for $ater molecules is more positi&e.

12O easier to reduce.

Electrolysis of dil&ted NaCl sol&tion

Anode

Cl2 ,(# & 2e- 2Cl- ,a%# E; < +?.

O2 ,(# & *& ,a%# & e- 2*2O ,l# E; < +?.2

,n dilute solution/ $ater $ill "e selected foroxidation "ecause of its lo$er Eo.



ctrolysis of Concentrated NaCl sol&ti



ctrolysis of Concentrated NaCl sol&ti

Cathode

2*2O ,l# & 2e- *2 ,(# & 2O*- ,a%# E; < -;.= +a

&

,a%# & e-

+a ,s# E; < -2.F?

E; for $ater molecules is more positi&e

12O easier to "e reduceAnode

Cl2 ,(# & 2e- 2Cl- ,a%# E; < +?.

O2 ,(# & *& ,a%# & e- 2*2O ,l# E; < +?.2

,n concentrated solution/ chloride ions $ill "eoxidised "ecause of its high concentration.

R i i l d



Reactions involved

2*2O ,l# & 2e

-

*2 ,(# & 2O*

-

,a%# E;

< -;.= 2Cl- ,a%# Cl2

,(# & 2e- E; < -?.

Cellreaction: 2*2O,l# & 2Cl- Cl2,(# & *2,(# & 2O*-,a%#

E;cell < -2.?A

Cathode8

Anode8



0a2'O aqueous solution contains 0a+ ion/ 'O2- ion

and $ater molecules

On electrol#sis/ the cathode attracts 0a+ ion and 12O molecules

the anode attracts 'O2- ion and 12O molecules

Exercise

6redict the electrol#sis reaction $hen0a2'O solution is electrol#sed using platinum electrodes.

Soltion

Cathode



Anode

E; for $ater molecules is less positi&e

12O easier to oxidise

02O12- ,a%# & 2e- 20O

2- ,a%# E; < +2.;?

O2 ,(# & *& ,a%# & e- 2*2O ,l# E; < +?.2

Cathode

2*2O ,l# & 2e- *2 ,(# & 2O*- ,a%# E; < -;.=

+a& ,a%# & e- +a ,s# E; < -2.F?

E; for $ater molecules is more positi&e12O easier to reduce

E ti



Cathode < 12 gas is produced and solution "ecome:a/ic at cathode "ecause O1- ions are formed

%node < O2 gas is produced and solution "ecomeacidic at anode "ecause 1+ ions are formed

E!ation

Cathode8

Anode8 E; < -?.2

2*2O ,l# & 2e- *2 ,(# & 2O*- ,a%# E; < -;.=

2*2O ,l# O2 ,(# & *& ,a%# & e-

Cell

Reaction:

E;cell < -2.; 2*2O,l# O2,(# & 2*2,(#

%araday.s /a, of



%araday s /a, ofElectrolysis

5escri"es the relationship "et$een the amount ofelectricit# passed through an electrol#tic cell andthe amount of su"stances produced at electrode.

%araday.s %irst /a,

ates that the !antity o" s#stance "ormed at a

ectrode is directly $ro$ortional to the !antityelectric char%e s$$lied&



%araday.s 1st /a,

m α Q

3 electric char(e in coulombs ,C#

m 3 mass of substance dischar(ed



The End



Q = It

3 electric char(e in coulombs ,C#

I 3 current in amperes ,A#

t 3 time in second ,s#

7arada# constant (7)is the charge on ? mole of electron

' ( ) *+ ,--C

Exam$l



Exam$l

e%n aqueous solution of Cu'O is electrol#sed using acurrent of ;.?>; % for > hours. Calculate the mass

of copper deposited at the cathode.

Answer

Electric charge/ G < Current (,) x time (t)

G < (;.?>; %) x ( > x ; x ; )s

G < 2F;; C

? mole of electron H ? 7 H A >;; C

0o. of e- passed through <2F;;

A >;;

< ;.;2= mol

C 2+ ( ) 2 C ( )



Cu2+ (aq) + 2e- → Cu (s)

7rom equation:

2 mol electrons → ? mol Cu

;.;2= mol electrons→ ;.;? mol Cu

@r for Cu < .>@ass of Copper deposited < ;.;? x .>

< ;.==A g

26772095 matriculation chemistry electrochemistry

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