Redox Reactions - Chemistry Notes Class 11 - JEE PDF Download

Oxidation number denotes theoxidation state of an element in a compound ascertained according to a setof rules formulated on the basis thatelectron in a covalent bond belongsentirely to more electronegative element.

Calculation of oxidation number-

1. O. S. of all the elements in their elemental form (in standard state) is taken as zero O. S. of elements in Cl₂, F₂, O₂, P₄, O₃,Fe(s), H₂, N₂, C(graphite) is zero.

2. Common O. S. of elements of group one (1st) is one. Common O. S. of elements of group two (2nd) is two.

3. For ions composed of only one atom, theoxidation number is equal to the chargeon the ion.

4. The oxidation number of oxygen in most compounds is –2.While in peroxides (e.g., H₂O₂, Na₂O₂), eachoxygen atom is assigned an oxidationnumber of –1, in superoxides (e.g., KO₂,RbO₂) each oxygen atom is assigned anoxidation number of –(½).

5. In oxygendifluoride (OF₂) and dioxygendifluoride (O₂F₂), the oxygen is assignedan oxidation number of +2 and +1,respectively.

6. The oxidation number of hydrogen is +1 but in metal hydride its oxidation no. is–1.

7. In all its compounds, fluorine has anoxidation number of –1.

8. The algebraic sum of the oxidation numberof all the atoms in a compound must bezero.

9. In polyatomic ion, the algebraic sumof all the oxidation numbers of atoms ofthe ion must equal the charge on the ion.

Stocknotation:the oxidation number is expressed by putting a Romannumeral representing the oxidation numberin parenthesis after the symbol of the metal inthe molecular formula. Thus aurous chlorideand auric chloride are written as Au(I)Cl and Au(III)Cl₃. Similarly, stannous chloride andstannic chloride are written as Sn(II)Cl₂ and Sn(IV)Cl_4.

Oxidation: An increase in the oxidationnumber

Reduction: A decrease in the oxidationnumber

Oxidising agent: A reagent which canincrease the oxidation number of an elementin a given substance. These reagents are calledas oxidants also.

Reducing agent: A reagent which lowers the oxidation number of an element in a givensubstance. These reagents are also called asreductants.

Redox reactions: Reactions which involvechange in oxidation number of the interactingspecies

Balancing of redox reactions:

Oxidation Number Method: Write the net ionic equation for the reaction of potassium dichromate(VI), K₂Cr₂O₇ with sodium sulphite,Na₂SO₃, in an acid solution to give chromium(III) ion and the sulphate ion

Step 3: Calculate the increase anddecrease of oxidation number, and make them equal:

Step 4: Balance the charge by adding H⁺ as the reaction occurs in theacidic medium,

Step 5: Balance the oxygen atom by adding water molecule

Half Reaction Method 

balance the equation showing the oxidation of Fe²⁺ ions to Fe³⁺ ions by dichromate ions (Cr₂O₇)^2– in acidic medium,wherein, Cr₂O₇^2– ions are reduced to Cr³⁺ ions.

Step 1: Produce unbalanced equation for thereaction in ionic form :

Fe²⁺(aq) + Cr₂O₇^2- (aq) → Fe^{3+ (}aq) + Cr³⁺(aq)

Step 2: Separate the equation into halfreactions:

                           +2 +3

Oxidation half : Fe²⁺ (aq) → Fe³⁺(aq)

                       +6 –2 +3

Reduction half :Cr₂O₇ ^2– (aq) → Cr³⁺(aq)

Step 3: Balance the atoms other than O andH in each half reaction individually.

Cr₂O₇ ^2– (aq) → Cr³⁺(aq)

Step 4: For reactions occurring in acidicmedium, add H₂O to balance O atoms and H ⁺ to balance H atoms.

Cr₂O₇ ^2– (aq) +14 H⁺→ Cr³⁺(aq) + 7H₂O (l)

Step 5: Add electrons to one side of the halfreaction to balance the charges. If need be,make the number of electrons equal in the twohalf reactions by multiplying one or both halfreactions by appropriate coefficients.

Fe²⁺ (aq) → Fe³⁺ (aq) + e^– 

Cr₂O₇ ^2–(aq) + 14H⁺ (aq) + 6e– → 2Cr³⁺(aq) +7H₂O (l)

6Fe²⁺ (aq) →6 Fe³⁺ (aq) +6 e^– 

Step 6: We add the two half reactions toachieve the overall reaction and cancel theelectrons on each side. This gives the net ionicequation as :

6Fe²⁺(aq) + Cr₂O₇ ^2– (aq) + 14H⁺ (aq) → 6 Fe³⁺(aq) +2Cr ³⁺(aq) + 7H₂O(l)

A redox couple is defined as havingtogether the oxidised and reduced forms of asubstance taking part in an oxidation orreduction half reaction. Represented as

Zn²⁺/Zn and Cu²⁺/Cu.

¤  Electrochemical cells are the devices which are used to get electric current by using chemical reaction.

The potential associated with eachelectrode is known as electrode potential. If the concentration of each species taking partin the electrode reaction is unity (if any gasappears in the electrode reaction, it is confinedto 1 atmospheric pressure) and further thereaction is carried out at 298K, then thepotential of each electrode is said to be the Standard Electrode Potential. 

• SHE is used to measure electrode potential and its standard electrode potential is taken as 0.00 V.

1. Redox reactions are those reactions in which oxidation and reduction takes place simultaneously

2. Classical view of redox reactions

• Oxidation is addition of oxygen / electronegative element to a substance or removal of hydrogen / electropositive element from a substance
• Reduction is removal of oxygen / electronegative element from a substance or addition of hydrogen / electropositive element to a substance

3. Redox reactions in terms of Electron transfer

• Oxidation is defined as loss of electrons by any species
• Reduction is defined as gain of electrons by any species

4. In oxidation reactions there is loss of electrons or increase in positive charge or decrease in negative charge

5. In reduction reactions there is gain of electrons or decrease in positive charge or increase in negative charge

6. Oxidising agents are species which gain one or more electrons and get reduced themselves

7. Reducing agents are the species which lose one or more electrons and gets oxidized themselves

8. Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules. These rules are formulated on the basis that electron in a covalent bond belongs entirely to the more electronegative element.

9. Rules for assigning oxidation number to an atom

• Oxidation number of Hydrogen is always +1 (except in hydrides, it is -1).
•  Oxidation number of oxygen in most of compounds is -2. In peroxides it is (- 1). In superoxides, it is (-1/2). In OF₂ oxidation number of oxygen is +2.In O₂F₂ oxidation number of oxygen is +1
• Oxidation number of Fluorine is -1 in all its compounds
• For neutral molecules sum of oxidation number of all atoms is equal to zero
• In the free or elementary state, the oxidation number of an atom is always zero. This is irrespective of its allotropic form
• For ions composed of only one atom, the oxidation number is equal t the charge on the ion
• The algebraic sum of the oxidation number of all the atoms in a compound must be zero
• For ions the sum of oxidation number is equal to the charge on the ion
• In a polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must be equal to the charge on the ion

10. Oxidation state and oxidation number are often used interchangeably

11. According to Stock notation the oxidation number is expressed by putting a Roman numeral representing the oxidation number in parenthesis after the symbol of the metal in the molecular formula

12.Types of Redox Reactions

• Combination Reactions: Chemical reactions in which two or more substances (elements or compounds) combine to form a single substance
• Decomposition Reactions: Chemical reactions in which a compound break up into two or more simple substances
• Displacement Reactions: Reaction in which one ion(or atom)in a compound is replaced by an ion(or atom) of other element
• a. Metal Displacement Reactions: Reactions in which a metal in a compound is displaced by another metal in the uncombined state
• b. Non-metal Displacement Reactions: Such reactions are mainly hydrogen displacement or oxygen displacement reactions
• Disproportionation Reactions: Reactions in which an element in one oxidation state is simultaneously oxidized and reduced

13.Steps involved in balancing a Redox reaction by oxidation number method

• Write the skeletal redox reaction for all reactants and products of the reaction
• Indicate the oxidation number of all the atoms in each compound above the symbol of element
• Identify the element/elements which undergo change in oxidation numbers
• Calculate the increase or decrease in oxidation number per atom
• Equate the increase in oxidation number with decrease in oxidation number on the reactant side by multiplying formula of oxidizing agent and reducing agents with suitable coefficients
• Balance the equation with respect to all other atoms except hydrogen and oxygen
• Finally balance hydrogen and oxygen. For balancing oxygen atoms add water molecules to the side deficient in it. Balancing of hydrogen atoms depend upon the medium
• a. For reactions taking place in acidic solutions add H ⁺ ions to the side deficient in hydrogen atoms
• b. For reactions taking place in basic solutions add H₂O molecules to the side deficient in hydrogen atoms and simultaneously add equal number of OHions on the other side of the equation
• Finally balance the equation by cancelling common species present on both sides of the equation

14.Steps involved in balancing a Redox by Ion-Electron Method(Half reaction method)

• Find the elements whose oxidation numbers are changed. Identify the substance that acts as an oxidizing agent and reducing agent
• Separate the complete equation into oxidation half reaction and reduction half reaction
• Balance the half equations by following steps
• i. Balance all atoms other than H and O
• ii. Calculate the oxidation number on both sides of equation .Add electrons to whichever side is necessary to make up the difference
• iii. Balance the half equation so that both sides get the same charge
• iv. Add water molecules to complete the balancing of the equation
• Add the two balanced equations. Multiply one or both half equations by suitable numbers so that on adding two equations the electrons are balanced

15. Application of Redox reactions: Redox Titrations

• Potassium permanganate in redox reactions: Potassium permanganate (KMnO₄) is very strong oxidizing agent and is used in determination of many reducing agents like Fe^2+, oxalate ions etc. It acts as self indicator in redox reactions. 

           Equation showing KMnO₄ as an oxidising agent in acidic medium is: 

          MnO₄  + 8H⁺ + 5e^- → Mn²⁺ + 4H₂O

• Acidified Potassium dichromate (K₂Cr₂O₇) in redox reactions: K₂Cr₂O₇ is used as an oxidizing agent in redox reactions. Titrations involving K₂Cr₂O₇ uses diphenylamine and potassium ferricyanide (external indicator).

           Equation showing K₂Cr₂O₇ as an oxidising agent in acidic medium is:

• Iodine (I₂) in redox reactions: I₂ acts as mild oxidising agent in solution according to equation

            I₂ + 2e^-→ 2I^-

16.Direct redox reaction: Redox reactions in which reduction and oxidation occurs in same solution (i.e. same reaction vessel).In these reactions transference of electrons is limited to very small distance.

17.Indirect redox reactions: Redox reactions in which oxidation and reduction reactions take place in different reactions vessels and thus transfer of electrons from one species to another does not take place directly

18.Electrochemical cell is a device that converts chemical energy produced in a redox reaction into electrical energy. These cells are also called Galvanic cells or Voltaic cells

19. The electrode at which oxidation occurs is called anode and is negatively charged

20.The electrode at which reduction takes place is called cathode and is positively charged

21.In an electrochemical cell the transfer of electrons takes place from anode to cathode

22. In an electrochemical cell the flow of current is from cathode to anode

23.In the electrochemical cell, the electrical circuit is completed with a salt bridge. Salt bridge also maintains the electrical neutrality of the two half cells

24. A salt bridge is a U shaped tube filled with solution of inert electrolyte like sodium chloride or sodium sulphate which will not interfere in the redox reaction. The ions are set in a gel or agar agar so that only ions flow when inverted

25.Electrical potential difference developed between the metal and its solution is called electrode potential. It can also be defined as tendency of an electrode in a half cell to gain or lose electrons

26.Oxidation potential is the tendency of an electrode to lose electrons or to get oxidized

27.Reduction potential is the tendency of an electrode to gain electrons or get reduced

28.In an electrochemical cell, by the present convention, the electrode potentials are represented as reduction potential

29.The electrode having a higher reduction potential will have a higher tendency to gain electrons

30. By convention, the standard electrode potential of hydrogen electrode is 0.00 volts

31. A redox couple is defined as having together oxidized and reduced forms of a substance taking part in an oxidation or reduction half reaction

32.The difference between the electrode potentials of eth two electrodes constituting the electrochemical cell is called EMF(Electromotive force) or the cell potential

33.. EMF = 

34. A negative E^θ means that the redox couple is a stronger reducing agent than the H⁺ /H₂ couple

35.A positive E^θ  means that the redox couple is a weaker reducing agent than the H ⁺ /H₂ couple

Questions: 

1. Define oxidation and reduction in terms of oxidation number.

Ans Increase in Oxidation Number is Oxidation and decrease in Oxidation Number is called reduction.

2. What is meant by disproportionation? Give one example.

Ans : In a disproportionation reaction an element simultaneously oxidized and reduced.

P₄ + 3OH^- +3H₂O→ PH₃ +3H₂PO^{2 -} 

3. What is O.N. of sulphur in H2SO4?

Ans: +6

4. Identify the central atom in the following and predict their O.S. HNO₃

Ans: central atom:- N; O.S. +5

5. Out of Zn and Cu which is more reactive?

Ans: Zn.

6. What is galvanization?

Ans: Coating of a less reactive metal with a more reactive metal e.g.- coating of iron surface with Zn to prevent rusting of iron.

7. How is standard cell potential calculated using standard electrode potential?

Ans: E ⁰ cell = E⁰ cathode – E ⁰ anode

8. What is O.S. of oxygen in H₂O₂?

Ans:  -1.

9. The formation of sodium chloride from gaseous sodium and gaseous chloride is a redox process justify.

Ans: Na atom get oxidize and Cl is reduced