1800-212-7858 (Toll Free)
9:00am - 8:00pm IST all days

or

Thanks, You will receive a call shortly.
Customer Support

You are very important to us

For any content/service related issues please contact on this toll free number

022-62211530

Mon to Sat - 11 AM to 8 PM

Redox Reactions And Electrochemistry

Share this:

Redox Reactions and Electrochemistry PDF Notes, Important Questions and Formulas

Redox Reactions and Electro-chemistry

Introduction to Oxidation & Reduction

Oxidation is the process involving addition of oxygen or removal of hydrogen or both.

Example of oxidation:

Fe   → FeO                      (Addition of Oxygen)

CH3CH2OH  →  CH3CHO (removal of hydrogen)                          

CH3CH2OH  →  CH3COOH   (both)

Reduction is just the reverse process. It involves addition of hydrogen or removal of oxygen or both.

Examples of reduction:

ZnO  →  Zn   (removal of oxygen)

C2H4  →  C2H(addition of hydrogen)

CH3COOH  →  CH3CH2OH  (both)

Limitations of old concept:

(i) It include only the reactions involving oxygen and hydrogen atoms: For example:

                    Fe → FeO and Fe → FeCl2

Both the processes are similar because in both the processes, iron is forming a compound from its elemental state, in which its valency is two. The first process is oxidation but second process is not oxidation but second process is not oxidation by this definition, because it does not involve oxygen or hydrogen atoms.

(ii) There may be some process even involving oxygen and hydrogen atoms, which cannot be classified as oxidation or reduction. For example:

               Fe → FeCl(Oxidation, addition of Cl atoms)

 

(Oxidation, addition of – OH groups)

The process involves addition of oxygen as well as hydrogen and hence it cannot be classified as oxidation reduction by this definition.

(iii) Conversation of sodium into sodium hydride is oxidation process, but it is reduction by his definition

Na → NaH      (addition of hydrogen)

Modification:

The old definition may be modified by replacing oxygen atom with electronegative atom or group and replacing hydrogen atom with electropositive atom or group.

   Fe → FeCl2    (Oxidation, addition of Cl atoms)

(Oxidation, addition of –OH group)

MODERN CONCEPT

1. In terms of electrons:

Oxidation is the process involving loss of electron and reduction is the process involving gain of electrons.

Example:

Zn → Zn2++2e             (Oxidation)

2Cl→ Cl2+2e                  (Oxidation)

Cu2++2e → cu                  (Reduction)

MnO4+ 8H++5e → Mn2++4H2O (Reduction)

Cl+12 OH-→ 2ClO3_+6H2O +10e (Oxidation)

The number of electrons lost or gained can be determined by first conserving the atoms and then, conserving the charge.

     2. In terms of oxidation state:

Oxidation is the process involving increase in oxidation state of one or more element while reduction is the process involving decrease in oxidation state of one or more element while reduction is the process involving decrease in oxidation state of one or more element. Oxidation state of any atom in any molecule or ion may be defined as arbitrary charge assigned to that atom according to some well-defined rules.

Rules for Determination of Oxidation State

If the bonded atoms are of element, distribute the shared electron equally between them and if the bonded atoms are different, count shared electron pair for more electronegative atom. Determine the net charge developed on the atoms after re-distribution of shared electrons. It will represent the oxidation states of the atoms. Examples

(i) HCL Molecule

1p    17p

H  ⎻  Cl

 

(ii) ClMolecule

17p      17p

Cl     ⎻   Cl

 

(iii) H2O2 Molecule:

+1-1-1+1

H-O-O-H

The re-distribution of electron may be made simply by assuming the covalent the covalent bonds,, ionic and assigning the charge on atoms on the basis of electro negativities and bond order.

Determination of Oxidation State

Rules I:

The oxidation state of any atom is its elemental state is zero.

Rules II:

The maximum oxidation state of any atom will be equal to (+group number) and minimum oxidation state will be equal to (group number –8), where group numbers are in roman numerals. For example, Sulphur (S) is member of group VI A and hence its maximum oxidation state is +6 and minimum is = (6–8) = –2

Rule III:

The sum of oxidation state of all the atoms in a molecule is zero and for ions, it is equal to the ionic charge.

Rule IV:

The oxidation states of some elements are fixed in all their compounds.

+1: Alkali metals (Li, Na, K, Rb, Cs, Fr) and Ag

+2: Alkaline earth metals (Be. Mg, Ca, Sr, Ba, Ra) and Zn

+3: Al –1: F

Rule V: 

Oxidation state of hydrogen is +1 in all of its compounds, except the metal hydrides, where it –1.

Rule VI:

Oxidation state oxygen is -2 in all of compounds except

  1. Peroxide like Na2O2, H2O2, BaO2, etc, where it is -1.
  2. Superoxides like KO2, RbO2, etc, where it is ½.
  3. Some other binary compounds of alkali metals and oxygen like KO3(O.S. of O=1/3), Rb2O3(O.S. of O=-2/3)etc
  4. Oxides of fluorine, where it is positive states. For example: O.S of O in OF2, O2F2, O3F2, etc are +2, +1,+2/3, respectively.

Rule VII:

The charges on different ions commonly used, should be known.

CO32

 

Carbonate ion

HCO3

Hydrogen carbonate ion

 

SiO44

Silicate ion

 

PO43

Phosphate ion

 

HPO42

Hydrogen phosphate ion

H2PO4

 

Dihydrogen phosphate ion

HPO32

Phosphite ion

 

NO32

Nitrate ion

 

NO2

Nitrate ion

 

SO42

Sulphate ion

 

S2

Sulphide ion

 

S2

Pyrite ion

 

S2O72

Disulphate ion

 

S2O32

Thiosulphate ion

 

S2O82

Peroxodisulphate ion

 

ClO

Hypochlorite ion

 

ClO3

Chlorate ion

 

ClO2

Chlorite ion

 

ClO4

Perchlorate ion

 

Rule VIII:

In the complex compound, the overall charge on ligand should be considered in place of considering the charges on individual atoms.

Neutral Ligands:

H2O, NH3, CO, NO, pyridine (Py), ethylenediamine (en), triphenylphoshine (Ph3P), etc

Negative Ligands:

X, OH, NH2, NO3, C2O4, (ox), O2,O22,SO42,S2O32,CNO⎻,etc

Positive Ligands:

 begin mathsize 12px style straight N with plus on top straight O comma text   end text stack NO subscript 2 with plus on top end style

Types of Reactions

Redox reactions: These are the reaction involving oxidation as well as reduction. Examples:

(i) begin mathsize 12px style straight Z with 0 on top straight n plus 2 straight H with plus 1 on top Cl rightwards arrow straight Z with plus 2 on top nCl subscript 2 plus straight H with 0 on top subscript 2 end style

(Oxidation: Zn, Reduction=HCl)

(ii) begin mathsize 12px style straight K straight M with plus 7 on top nO subscript 4 plus straight H subscript 2 SO subscript 4 plus straight H subscript 2 straight S with negative 2 on top rightwards arrow straight K subscript 2 SO subscript 4 plus straight M with 2 plus on top nSO subscript 4 plus straight H subscript 2 straight O plus straight S with 0 on top end style

(Oxidation: H2S, reduction=KMnO4)


Electrochemical cell and nearest equation Electrochemical cells

An electrochemical cell consists of two electrodes (metallic conductors) in contact with an electrolyte (an ionic conductor). An electrode and its electrolyte comprise an 

Electrode Compartment.

Electrochemical Cells can be classified as: 

  1. Electrolytic Cells in which a non-spontaneous reaction is driven by an external source of current.
  2. Galvanic Cells which produce electricity as a result of a spontaneous cell reaction.

Note: In a galvanic cell, cathode is positive with respect to anode. In an electrolytic cell, anode is made positive with respect to cathode.

GALVANIC CELL

This cell converts chemical energy into electrical energy. Galvanic cell is made up of two half cells i.e., anodic and catholic. The cell reaction is of redox kind. Oxidation takes place at anode and reduction at cathode. It is also known as voltaic cell. It may be represented as shown in Fig. Zinc rod immersed in ZnSO4 behaves as anode and copper rod immersed in CuSO4 behaves as cathode.

 

Oxidation takes place at anode.

Zn→Zn2++2e- (loss of electron: oxidation)

Reduction takes place at cathode:

Cu2++2e-→Cu (gain of electron; reduction)

Overall Process: Zn(s)+Cu2+→Cu(s)+Zn2+ In galvanic cell like Daniell cell: electron flow from anode (zinc rod) to the cathode (copper rod) through external circuit; zinc dissolves as Zn2+ ;Cu2+ ion in the cathode cell picks up two electron and become deposited at cathode.

REPRESENTATION OF A CELL (IUPAC CONVENTIONS)

Let us illustrate the convention taking the example of Deniel cell

  1. Anodic half cell is written on left and cathodic half cell on right hand side.                                                                            Zn(s) ∣ZnSO4(sol)∥ CuSO4(sol)∣Cu(s)
  2. Two half cell are separated by double vertical lines: Double vertical lines indicate slat bridge or any type of porous partition.
  3. EMF (electromotive force) may be written on the right hand side of the cell.
  4. Single vertical lines indicate the phase separation between electrode and electrolyte solution.                                    Zn∣Zn2+∥Cu2+∣Cu
  5. Invert electrode are separated in the bracket                                                                                                                Zn∣ZnSO4∥H+∣H2,Pt

RELATIONSHIP BETWEEN G AND ELECTRODE POTENTIAL

Let n, faraday charge is taken out from a cell of e.m.f (E) then electrical work done by the cell may be calculated as,

Work done=Charge × Potential =nFE

From thermodynamic we know that decrease in Gibbs free energy of a system is a measure of Gibbs free energy of a system is a measure of reversible or maximum obtainable work by the system if there is no work due to volume expansion

∴ △G=⎻nFE

Under standard state △G0=⎻nEF0   …..(1)

From thermodynamic we know, △G= negative for spontaneous process. Thus from eq.

  1. it is clear that the EMF should be + ve for a cell process to feasible or spontaneous.
  2. When △G = positive, E= negative and the cell process will be non spontaneous.

Reactions                     G               E

Spontaneous                  (-)              (+)

Non-Spontaneous          (+)              (-)

Equilibrium                     0                0

Standard free energy change of a cell may be calculated by electrode Potential data.

Substituting the value of E0 (i.e. standard reduction potential of   

cathode-standard reduction potential of anode) in eq.(i) We may get  △G.

Concept of electromotive porce (EMF) of A cell

Electron flows from anode to cathode in external circuit due to a pushing effect called or electromotive force (e.m.f). EMF is called as cell potential. Unit of e.m.f. of cell is volt.

EMF of cell may be calculated as:

Ecell=reduction potential of cathode – Reduction potential of anode

Similarly, standard e.m.f. of the cell (E0) may be calculated as

E0cell=Standard reduction potential of cathode – Standard reduction potential of anode.

DIFFERENT TYPES OF HALF-CELLS AND THEIR REDUCTION POTENTIAL

(1) Gas-Ion Half Cell:

In such a half cell, an inert collector of electrons, platinum or graphite is in contact with gas and a solution containing a specified ion. One of the most important gas-ion half cell is the hydrogen–gas-hydrogen ion half cell. In this cell, purified H2 gas at a constant pressure is passed over a platinum electrode which is in contact with an acid solution.

H+(aq) + e⇌ 1/2 H2



<!DOCTYPE html>
<html>
<head><meta charset="UTF-8"><title>Error 500</title></head>
<body>
<h1>Error 500</h1>

<!--

-->



</body>
</html>

(2) Metal-Metal Ion Half Cell:

This type of cell consist of a metal M is contact with a solution containing Mn+ ions.

Mn+(aq) + ne⇌ M(s)



<!DOCTYPE html>
<html>
<head><meta charset="UTF-8"><title>Error 500</title></head>
<body>
<h1>Error 500</h1>

<!--

-->



</body>
</html>

(3) Metal-Insoluble Salt-Anion Half Cell:

In this half cell, a metal coated with its insoluble.

CO32

 

Carbonate ion

HCO3

Hydrogen carbonate ion

 

SiO44

Silicate ion

 

PO43

Phosphate ion

 

HPO42

Hydrogen phosphate ion

H2PO4

 

Dihydrogen phosphate ion

HPO32

Phosphite ion

 

NO32

Nitrate ion

 

NO2

Nitrate ion

 

SO42

Sulphate ion

 

S2

Sulphide ion

 

S2

Pyrite ion

 

S2O72

Disulphate ion

 

S2O32

Thiosulphate ion

 

S2O82

Peroxodisulphate ion

 

ClO

Hypochlorite ion

 

ClO3

Chlorate ion

 

ClO2

Chlorite ion

 

ClO4

Perchlorate ion

 
Show more

NEET Tests & Papers Solutions

Chemistry syllabus

Purchase Our Experts Course Packages

Enroll now to crack NEET

Testimonials

Ask Experts for NEET

Queries asked on Sunday and after 7 pm from Monday to Saturday will be answered after 12 pm the next working day.

Chat with us on WhatsApp