• Oxidation and Reduction • Oxidation - Oxidation may be defind as “ A process which involves the addition of oxygen or any other electronegative element, or as a process which involves the removal of hydrogen or any electropositive element.” • For example- (i) 2Mg (s) + O 2(g) ⟶ 2Mg O (Addition of oxygen) (ii) Mg (s) + Cl 2 (g) ⟶ MgCl 2 (Addition of electronegative element chlorine) (iii) 2H 2 S (g) + O 2 (g) ⟶ 2S (s) +2H 2 O (Removal of hydrogen) (iv) 2KI (aq) + H 2 O (l) + O 3(g) ⟶ 2KOH (aq) + I 2(s) + O 2 (Removal of electropositive element,potassium)
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• Oxidation and Reduction
• Oxidation -
Oxidation may be defind as “ A process which involves the addition of oxygen
or any other electronegative element, or as a process which involves the removal of hydrogen
or any electropositive element.”
• For example-
(i) 2Mg(s) + O2(g) ⟶ 2Mg O (Addition of oxygen)
(ii) Mg(s) + Cl2 (g) ⟶ MgCl2 (Addition of electronegative element chlorine)
(iii) 2H2S (g) + O2 (g) ⟶ 2S (s) +2H2O (Removal of hydrogen)
▪ Oxidising agent (oxidant) and Reducing agent (Reductant)-
Oxidising agent (oxidant)-
▪ It oxidises some other substance and itself is reduced.
▪ It loses oxygen or any electronegative element.
▪ It accepts hydrogen or any electropositive element.
▪ It gains one or more electrons.
▪ O.N. of effective element decreases.
▪ Reducing Agent (Reductants)-▪ It reduces some other species and itself is oxidised.
▪ It accepts oxygen or any other electronegative element.
▪ It loses hydrogen or any other electropositive element.
▪ It lose one or more electron.
▪ O.N.of effective element increases.
▪ The different terms used in decreasing oxidation-reduction process are given……..
TermIn terms Of Oxidation number change In terms of electron
transfer
Oxidation Increase Loss of electron
Reduction Decreases Gain of electron
Oxidising agent Decreases Gain of electron
Reducing agent Increases Supplies electrons
Substance oxidised Increases Loss of electrons
Substance reduced Decreases Gains electrons
• Metal tendency⟶Oxidised
(Oxidation) M⟶M+ +ne-
“The potential difference between the electrode(metal) and its ion in electrolyte (under equilibrium) is electrode potential or Oxidation potential.”
Metal tendency-----------Reduced(reduction) Mn+ + ne- ⟶ M
from solution from metal
“The potential difference between the electrode(metal) and its ion in electrolyte (under equilibrium) is electrode potential or Reduction potential.”
• Eelectrode Potential -
When a metal rod is placed in
asolution of its own ions, the rod becomes more electropositive orelectronegative with respect to the solution.Thus, a potentialdifference is established between the metal and the solution. Thisdifference of potential is known as electrode potential.It is denotedby ‘E’.
“OR”
“ The tendency of an electrode to lose or gain electrons when it is incontact with its own ions in solution, is called electrode potential.”
Since the tendency to gain electrons means also the tendency to getreduced, this tendency is called reduction potential.
• Similarly ,the tendency to lose electrons means the tendency to get oxidised.Hence ,this tendency is called oxidation potential.
• Example - When zinc rod is placed in the solution ZnSO4 , zinc rod becomeselectronegative and a potential difference is established between the metal and zinc ions.Thepotential difference so produced is called the potential of zinc electrode.
• Factors affecting the Electrode Potential-
1-Nature of metal and its ion-
Zn⟶Zn2+
Fe⟶Fe2+
Fe⟶Fe3+
2-Concentration of ions-
In case of solution concentrarion term is defined in term of molar while incase of gas concentration is defined interm of partial pressure.
3- Temperature- (25+273)= 298 K (250c)
• Standard electrode potential-
concentration [Mn+]= 1m
Presuure in case of gas= 1atm/1bar
Temperature= 298K (250C)
• The standard electrode potential can be defined as the electrode potential when all the reactants and products are at their unit concentration at 250C and 1 atm (atmospheric pressure), It is represented by E0.
• Standard electrode potential= SEP (already given)
• Oxidation M⟶Mn+ + ne-
Em/mn+⟶ oxidation potential
E0m/m
n+ = standard oxidation potential
Reduction Xn+ + ne- ⟶ X
• Exn+/x = reduction potential
• E0xn+/x = standard electrode potential
According to IUPAC
Zn⟶Zn2++2e-
E0Zn/Zn
2+ =+0.76 v (SOP)
Zn2++2e- ⟶Zn
E0Zn
2+/Zn = -0.76 v (SRP)
• Acoording to IUPAC
SRP⟶SEP or Reduction potential = Electrode potential
Standard Reduction Potential = -Standard oxidation potential
• Standard electrode potential-
concentration [Mn+]= 1m
Presuure in case of gas= 1atm/1bar
Temperature= 298K (250C)
• The standard electrode potential can be defined as the electrode potential when all the reactants and products are at their unit concentration at 250C and 1 atm (atmospheric pressure), It is represented by E0.
EMF cell = ERP - ERP
ECell = E cathode - E anode
(RP) (RP)
OR
• Ecell = E right - E left
• (RP) (RP)
• ECell can calculate at any temperature or any
concentration by NERNST equation.
• E0Cell = Standard EMF of cell at Temperatur is 250c,
concentration is 1 molar and partial pressure is 1 atm. then
• E0Cell = E 0cathode - E0
anode
(RP) (RP)
• E 0cell = E0right - E0
left
• Question-
• Ni/Ni2+ (1M) II Ag+ (1M)/Ag
Left-anode Right-cathode
oxidation Reduction
Given: E0Ni/Ni
2+ = -0.25 v or E0Ni
2+/Ni = -(-0.25) = +0.25 v (reduction potential)
E0Ag
+/Ag = 0.80 v
calculate standard electrode potential.
• Solution:
E0Cell = E0
cathode - E0anode
(RP) (RP)= 0.80-(-025)= 0.80+0.25= 1.05 v
• Question: Calculate standard cell potential of DaniellCell-
Zn/Zn2+ (1M) II Cu2+ (1M)/Cu
• Given: E0Zn
2+/Zn= -0.76 v (RP)
E0Cu
2+/Cu = +0.34 v (RP)
Solution: E0Cell = E0
cathode - E0anode
= 0.34-(-0.76)
= 0.34+0.76
= 1.10 v
• Caculate standard cell potential
Zn/Zn2+ (1M) II Pb2+ (1M)/Pb
Given: E0Zn/Zn
2+ = +0.76 v
E0Pb
2+ /Pb = +0.76 v
E0Zn/Zn
2+ = -0.76 (RP)
E0 cell = E0 cathode - E0 anode
= +0.76 – (-0.76)
=+ 1.52 v
Effect of Metal Ion Concentration on the Magnitude of Electrode Potential –(Nerst Equation)-
We know that the standard oxidation potential of Zn/Zn2+
electrode in which the Zn rod is dipped in 1M solution of ZnSO4 is +0.76 volt.
If the concentration of the solution is not 1M, then the Oxidation potential
corresponding to the reaction:
Zn⟶ Zn2+ +2e-
Is not equal to +0.76 volt but is givenby:
E= E0zn / zn
2+ - log
Where n= No.of electrons in the electrode reaction =2
• Now since the concentration of the elements and solid =1, [Zn] = 1 and hence:
E= E0zn / zn
2+ - log [Zn +]
In general :
E= E0 - log [Zn +]
This is Nerst equation.
This is the equation for a half-cell oxidation reaction and hence:
Eox = E0– log [ion]
The value of Ered for the half –cell reduction is given by reversing the sign of Eox , i.e.
• Ered = - Eox= -{ E0- log [Ion]}
Ered = - E0 + log [Ion]
Effect Of metal Ion concentration on the Magnitude of e.m.f. of a given cell-
We have seen that the standard e.m.f. of Daniellcell represented as:
Zn/Zn2+ (1M) II Cu2+ (1M)/Cu
• In which the concentration of each of the metal ions is unity is give
E0Cell = E0
cathode - E0anode
= 0.34-(-0.76)
= 0.34+0.76
= 1.10volt
If the concentration of each of the metal ions is different as in the cell:
Zn/Zn2+ (0.1M) II cu2+ (0.1M)/ cu
The e.m.f. is given by:
Ecell = E0cell - log [ ]
Where n = No.of electrons involved in the redox reaction
Zn⟶Zn2+(O.1M) + 2e (oxidation)
Reduction half-reaction (cathode reaction)
Cu2+(0.01M)+2e ⟶ Cu
Cell reaction is represented as Zn+ Cu2+(0.01M) ⟶ Zn2+(0.1M) +Cu (n=2)
=
Now since concentration of the elements and solids =1, [Cu] or [Zn] and hence
= =
ECell = E0cell - log
= 1.10-0.0295x1
= 1.0705 volt.
Electro chemical series or e.m.f. Series –
“ The list of elements
or ions arranged in the decreasing order of their standard reduction
potential values is called electro-chemical or e.m.f. series or activity
series.
❖ From this series the following points should keep in mind while studying
the use of this series.
❖ The forword reaction shown in the central column of the series is a
reduction reaction and hence the standard electrode potential values are
standard reduction potential value.
❖ The species shown at the left hand side act as oxidising agents while those
given at the right hand side of the reduction reaction act as reducing agents.
❖ The standard Electrode potential values are decreasing from top to bottom
in the series i.e. possitive values (e.g. (F2/2F-, E0=+2.85 V e.g. (F2/2F-,
• e.g. (Li+/Li, E0=-3.04 V). Thus the species lying above hydroge have positive E0values while those lying below hydrogen have negative E0 .
• From top to bottom the oxidising power of oxidising agents is decrease (i.e. F2 to Li+
ion ) while the reducing power of reducing agents is increase in the same direction (i.e.from F- ion to Li metal). Thus F2 is the strongest oxidising agent while Li+ ion is theweakest oxidising agent. similarly F- ion is the weakest reducing agent while Li metal isthe strongest reducing agent.
• Aplication Of Electrochemical series:
• To compare the oxising power and reducing power of metals or non-metals-
The magnitude of E0 value of ametal or a non-metal gives a measure of itstendency to gain electrons to get reduced and hence to act as an oxidising agent or loseelectrons to get oxidised and hence to act as a reducing agent.
Gain of electrons (reduction)
M + e- ⟶ M-
oxidising agent Reducing agent
Loss of electrons (oxidation)
M ⟶ M+ + e-
Reducing agent oxidising agent
Thus the oxidising or reducing property of a given element in solution is measured by the magnitude of its E0 value.
Oxidising agent:
The species having +E0 i.e. ,the species lying above hydrogen in theelectrochemical series show a strong tendency to gain electrons to undergo reduction andhence they are strong oxidiasing agents.
As already said,with the decrease of +E0 value of a given species, the tendency of thatspecies to gain electrons to undergo reduction decrease and hence the oxidising power ofthat species also decrease.
• Examples: The metal like Au, Pt, Hg, Cu etc. which have positive E0 values are strong oxidising agents. The oxidising power of these metals is in the order :
Au >Pt > Hg > Ag> Cu, since Au has the maximum E0 value while Cu has the minimum E0 value as shown below:
Au3+ + 3e- ⟶ Au, E0= +1.42 V
Pt2+ + 2e- ⟶ Pt, E0= +1.20 V E0
values decrease Oxidising
power decrease
Hg2+ + 2e- ⟶ Hg, E0= +0.85 V
Ag+ + e - ⟶ Ag, E0= +0.80 V
Cu2 + 2e- ⟶ Cu , E0= +0.34 V
• (ii)- The non-metals like F2, Cl2, Br2,I2 and O2 which have +E0 values are all strong agents. F2 which has the maximum E0 value is the strongest oxidising agent while O2
which has the lowest E0 value is the weakest oxidising agent.Thus the oxidising power of these non-metals is in the order:
F2 > Cl2 > Br2 >I2 > O2
Since E0 values are also decreasing in the same trend as shown below:
• As per electrochemical series, E0 values decrease from F2 to Li+ ion , the tendency
of these species to gain electrons to undergo reduction decreases and hence their
oxidising power decreases from F2 to Li+ ion . F2 is the strongest oxidising agent
while Li+ ion is the weakest oxidising agent.
Reducing agent- The species having - E0 values i.e. the species lying below
hydrogen in the electrochemical series they are strong reducing agent. As already
said, with the decrease of E0 of a given species (i.e. as E0 values becomes more
negative), the tendency of that species to lose electrons to undergo oxidation
increases and hence the reducing power of that species also increase.
Examples: (i)- To apply the above rule we can consider the reducing property of alkali
metals (group IA) and alkaline earth metals (group IIA). E0 values of these metals
are given below:
Alkali metals (IA group)
Li+ + e- ⟶ Li, E0= -3.04V
Na+ + e- ⟶ Na, E0= -2.71V
K+ + e- ⟶ K, E0= -2.92V
Rb++ e- ⟶ Rb, E0= -2.92V
Cs+ + e- ⟶ Cs , E0= -2.92V
E0 values of alkali metals show that these metals have high negative values and hence are strong reducing agent .Since Li has the minimum E0 value (i.e. Li has maximum negative E0 value) It is the strongest reducing agent of all the alkali metals.
Alkali earth metals (IIA group)
Be++ + 2e- ⟶ Be, E0= -1.70V
Mg+++ 2e- ⟶ Mg, E0= -2.38V
Ca+++ 2e- ⟶ Ca, E0= -2.76V
Sr++ + 2e- ⟶ Sr , E0= -2.89V
Ba++ + 2e- ⟶ Ba, E0= -2.92V
Similarly all the alkaline earth metal have high negative E0 values and hence ,likealkali metals behaves as a strong reducing agent.Since E0 values are decreasingfrom Be to Ba, the reducing power of these metals is also increasing in the samedirection, i.e. Be ,with maximun E0 is the weakest reducing agent while Ba whichhas minimum E0 value is the strongest reducing agent.
In electrochemical series from top to bottom F- ion is the weakest reducing agentwhile Li, metal is the strongest reducing agent.
Electropositive Character of Metals- Metals like K , Ca ,Na etc. which lieat the bottom of the series readily lose their outer-most shell electrons toform metal cations and hence are strongly electropositive (i.e. weaklyelectronegative) those lying at the top are not able to lose the outer-mostshell electrons to form the cations and hence are weakly electropositive(strongly electronegative).
• To predict whether a given metal will displace another metal from aqueous solution of its salt - A metal with lower standard reduction potential will displace another metal with higher standard reduction potential from the aqueous solution of its salt and the metal with higher standard reduction potential gets precipitated. This means that a metal will displace another metal from the aqueous Solution of its salt that lies above it in the electrochemical series.
• Examples: (i)- If a piece of Fe is placed in a solution of CuSO4 some of Fe goes in to solution as Fe++ (aq.) ions and Cu metal gets precipitated , i.e. Fe displaces Cu from CuSO4 solution as shown below:
Fe + CuSO4 ⟶ FeSO4 + Cu
Fe (s) + Cu2+ (aq.) ⟶ Fe2+ (aq.) + Cu(s)
• The displacement of Cu from CuSO4 solution by Fe is because of the fact
that E0Fe
2+/Fe = ( -0.44v ) is less than E0
Cu2+
/Cu = (+0.34V) Why the
displacement of Cu from CuSO4 solution by Fe is possible or why the
• Since the reducing agent viz, Zn lies above the oxidising agent Mg2+ in the electrochemical series, the reaction is not possible to occur in the direction shown.However ,the reverse reaction takes place spontaneously because the reducing agent viz. Mg lies below the oxidising agent namely Zn2+ ions in the electrochemical series.
Zn (s) + Fe2+ (aq.) ⟶ Zn2+ (aq.) + Fe(s) is predictable ,since E0Zn
2+/Zn (=-
0.76V) is less than E0Fe
2+/Fe =(-0.44).
This reaction is used for galvanising Fe to Zn to prevent rusting of Fe. Zn coating on Fe prevents the oxidation of Fe to Fe2+ by air . If the galvanised Fe is scratched and some iron is oxidised to Fe2+ by air , Fe2+ thus produced is immadiately reduced by Zn to Fe and thus the rusting of Fe is prevented.