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reduction

oxidation

CBSE Class 11 Chemistry Chapter 8 Revision Notes

Redox reactions are oxidation-reduction chemical processes in which the oxidation states of the reactants change. The term'redox' refers to the reduction-oxidation process. All redox reactions may be divided down into two types of reactions: reduction and oxidation.

Chapter 8: Redox Reactions Revision Notes

Oxidation

  • The addition of oxygen/electronegative element to a material or the removal of hydrogen/electropositive element from a substance is referred to as oxidation.

Reduction

  • The removal of oxygen or an electronegative element from a material, or the addition of hydrogen or an electropositive element to a substance, is referred to as reduction.

ELECTRON TRANSFER REACTION

  • Every redox reaction, according to the electronic notion, is made up of two processes known as half reactions.
  • Oxidation reaction: Oxidation reactions are half processes that entail the loss of electrons.
  • Reduction reaction: Reduction reactions are half reactions that include the gain of electrons.
  • Electron acceptor: An oxidising agent is a substance that accepts electrons.
  • Reducing agent: An electron donor.

Competitive Electron Transfer Reactions

Place a strip of metallic zinc in an aqueous solution of copper nitrate. Following changes will be seen after one hour.

(i) Strips are covered in a reddish metallic copper coating.

(ii) The solution's blue colour fades away.

(iii) When hydrogen sulphide gas is passed through the solution, white ZnS appears, as observed when the solution is made alkaline with ammonia.

Oxidation Number

  • It is the charge assigned to an atom of a compound that is equal to the number of electrons in that atom's valence shell that are gained or lost completely or to a large extent while that atom forms a bond in a compound.

Oxidation Number Assignment Rules

(i) In its simplest form, an element's oxidation number is zero.

The oxidation number of H2, 02, N2, and other gases is zero.

(ii) The oxidation number of a single monoatomic ion is equal to the ion's charge. For example, the oxidation number of the Na+ ion is +1, while the oxidation number of the Mg2+ ion is +2.

(iii) Oxygen's compounds have an oxidation number of -2. There are, however, certain exceptions.

Peroxides are a kind of compound. Na202, H202 \soxidation number of oxygen = – 1 In OF2 \sO.N. of oxygen = +2 02F2 \sO.N. of oxygen = +1

(iv) The oxidation number of hydrogen is + 1 in nonmetallic hydrogen compounds such as HCl, H2S, and H2O, but -1 in metal hydrides such as LiH, NaH, and CaH2.

(v) Metals have a positive oxidation number in compounds, whereas non-metals have a negative oxidation number. For example, in NaCl, the oxidation number of Na is +1, while the oxidation number of chlorine is -1.

(vi) When two non-metallic atoms are present in a compound, the atoms with the highest electronegativity are assigned a negative oxidation number, while the other atoms are assigned a positive oxidation number.

(vii) The algebraic total of all atoms in a compound's oxidation number equals zero.

(viii) The net charge on a polyatomic ion is equal to the sum of the oxidation numbers of all the atoms in the ion.

For instance, the sum of carbon atoms and three oxygen atoms is -2 in (C03)2.

Fluorine (F2) is such a highly reactive non-metal that it robs water of oxygen.

BALANCING REDOX REACTIONS

(i) The Method of Oxidation Numbers

(a) For each reactant and product, write the correct formula.

(b) The oxidation number change can be recognised by assigning the oxidation number.

(c) With regard to the reactants, calculate the increase and reduction in oxidation number per atom. If there are more than one atoms, multiply by the appropriate coefficient.

(d) Correct the equation in terms of all atoms. Hydrogen and oxygen atoms must also be balanced.

(e) Use H+ ions in the equation if the reaction is carried out in an acidic media. Use OH– ions if the medium is basic.

(f) By adding (H20) molecules to the reactants or products, the expression's hydrogen atoms can be balanced.

The balanced redox reaction occurs when both sides of the equation have the equal amount of oxygen atoms.

(ii) The Method of Half-Reaction

This approach balances two half equations individually before adding them together to create a balanced equation.

TITRATIONS

  • In these titrations, potassium permanganate (pink in colour) functions as an oxidising agent in the acidic media, whereas oxalic acid or various ferrous salts act as reducing agents.

  • These are some redox titration examples.

    Potassium permanganate serves as an indicator in each of these titrations. It's also known as a self-indicator. The end points are represented by the presence of pink colour in the solution.

    Potassium Dichromate Titration: In the presence of dil. H2S04, potassium dichromate can be used instead of potassium permanganate. The ionic equation for the FeS04 (Fe2+ ions) redox process is presented.

REDOX REACTIONS AND ELECTRODE PROCESSES

Electrochemical Cells

  • Electrochemical cells are devices in which the redox reaction is carried out indirectly and the energy loss manifests as electrical energy.

Electrolytic Cell

  • The cell that converts electrical energy to chemical energy. When a lead storage battery is recharged, for example, it becomes an electrolytic cell.
  • When a zinc rod is dipped in a copper sulphate solution, a redox reaction occurs, resulting in zinc being oxidised to Zn2+ ions and Cu2+ ions being reduced to metal.

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