Redox Reactions Class 11 Notes Chemistry Chapter 7 - CBSE

Chapter : 7

What Are Redox Reactions ?

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    Oxidation is defined as the addition of oxygen/electronegative element to a substance or removal of hydrogen or electropositive element from a substance.

    $$\text{2Mg(s) + O(g)}_{2}(g)\xrightarrow{}\\\text{2MgO(s)}\\\text{Mg(s) + Cl}_{2}(g)\xrightarrow{}\\\text{MgCl}_2\text{(s)}$$


    Reduction is defined as the removal of oxygen or electronegative element from a substance or addition of hydrogen or electropositive element to a substance.

    $$\text{2FeCl}_{3}(aq) + \text{H}_{2}(g)\xrightarrow{}\\\text{2 FeCl}_{2}(\text{aq}) + \text{2HCl(aq)}\\\text{2HgO(s)}\xrightarrow{}\\\text{2Hg(l) + O}_{2}(g)$$

    Oxidation Number Or Oxidation State

    It denotes the oxidation state of an element in a compound ascertained according to a set of rules formulated on the basis that electron pair in a covalent bond belongs entirely to more electronegative element. The idea of oxidation number has been invariably applied to define the following:

    • Oxidation: An increase in the oxidation number of the element in the given substance.
    • Reduction: A decrease in the oxidation number of the element in the given substance.
      • Oxidising Agent: A reagent which can increase the oxidation number of an element in a given substance. These reagents are called as oxidants also.
      • Reducing Agent: A reagent which lowers the oxidation number of an element in a given substance. These reagents are also called as reductants.
    • Redox reactions: Reactions which involve change in oxidation number of the interacting species.

    Rules For Assigning Oxidation Numbers

    • The oxidation number of an element in its free or uncombined state is zero. For example, H2 , 02, N2 etc. have oxidation number equal to zero.
    • In a single monoatomic ion, the oxidation number is equal to the charge on the ion. For example, Na+ ion has oxidation number of +1 and Mg2+ ion has +2.
    • Oxygen has oxidation number -2 in its compounds. However, there are some exceptions in the case of: Compounds such as peroxides. Na2 02 , H2 02, each oxygen atom is assigned an oxidation number of –1,
      Compounds such as superoxides (e.g., KO2 , RbO2), each oxygen atom is assigned an oxidation number of –(½). Another exception appears rarely, i.e. when oxygen is bonded to fluorine. In such compounds e.g., oxygen difluoride (OF2 ) and dioxygen difluoride (O2 F2), the oxygen is assigned an oxidation number of +2 and +1, respectively.
    • In non-metallic compounds of hydrogen like HCl, H2S, H2O oxidation number of hydrogen = + 1 but in metal hydrides oxidation number of hydrogen = -1 [LiH, NaH, CaH2 etc.]
    • In compounds of metals and non-metals metals have positive oxidation number while non-metals have
      negative oxidation number. For example, In NaCl. Na has +1 oxidation number while chlorine has -1.
    • If in a compound there are two non-metallic atoms the atoms with high electronegativity is assigned negative oxidation number while other atoms have positive oxidation number.
    • The algebraic sum of the oxidation number of all atoms in a compound is equal to zero.
    • In polyatomic ion the sum of the oxidation number of all the atoms in the ion is equal to the net charge on
      the ion.

    Types Of Redox Reactions

    Combination Reactions

    It involves the combination of two compounds to form a single compound in the form 

    $$\text{A + B}\xrightarrow{}\text{AB}.$$

    For example:

    $$\centerdot\space\text{H}_{2} + \text{Cl}_{2}\xrightarrow{}\text{2HCl}\\\centerdot\space\text{C + O}_{2}\xrightarrow{}\text{CO}_{2}\\\centerdot\space 4\text{Fe} + 3\text{O}_{2}\xrightarrow{}2\text{Fe}_{2}\text{O}_{3}$$

    Decomposition Reactions

    It involves the breakdown of a compound into different compounds. These reactions result in the breakdown of smaller chemical compounds in the form

    $$\text{AB}\xrightarrow{}\text{A + B}.$$

    For example:

    $$\centerdot\space\text{2NaH}\xrightarrow{}\text{2Na + H}_{2}\\\centerdot\space\text{2H}_{2}\text{O}\xrightarrow{}\text{2H}_{2}+\text{O}_{2}\\\centerdot\space\text{Na}_{2}\text{CO}_{3}\xrightarrow{}\text{Na}_{2}\text{O} +\text{CO}_{2} $$

    Displacement Reactions

    In this reaction an atom or an ion in a compound is replaced by an atom or an ion of another element. It can be represented in the form of

    $$\text{X + YZ}\xrightarrow{}\text{XZ + Y}.$$

    These reactions can be further categorised into :

    • Metal Displacement: In this
      type of reaction, a metal present in the compound is displaced by another metal. These types of reactions find
      their application in metallurgical processes where pure metals are obtained from their ores.
    • Non-metal Displacement:
      This type of reaction involves hydrogen displacement and sometimes rarely occurring reactions involving oxygen displacement.

    Disproportionat Reactions

    In this reaction a single reactant is oxidised and reduced.

    $$\text{P}_{4} + 3\text{NaOH} + 3\text{H}_{2}\text{O}\xrightarrow{}\\\text{3 NaH}_{2}\text{PO}_{2} +\text{PH}_{3}$$

    Balancing Of Redox Reactions

    Two methods are used to balance chemical equations for redox processes.

    Oxidation Number Method

    This method is based on the change in the oxidation number of reducing agent and the oxidising agent. Following steps are involved:

    • Step 1: Write the correct formula for each reactant and product.
    • Step 2: By assigning the oxidation change in oxidation number can be identified.
    • Step 3: Calculate the increase and decrease in oxidation number per atom with respect to the reactants. If
      more than one atom is present then multiply by suitable coefficient.
    • Step 4: Balance the equation with respect to all atoms. Balance hydrogen and oxygen atoms also.
    • Step 5: If the reaction is carried out in acidic medium, use H+ ions in the equation. If it is in basic medium use OH ions.
    • Step 6: Hydrogen atoms in the expression can be balanced by adding (H2O) molecules to the reactants or products. If there are same number of oxygen atoms on both sides of equation then it represents the balanced redox reaction.

    Half Reaction Method

    In this method two half equation are balanced separately and than added together to give balanced equation. Following steps are involved:

    • Step 1: Produce unbalanced equation for the reaction in ionic form.
    • Step 2: Separate the equation into half reactions.
    • Step 3: Balance the atoms other than O and H in each half reaction individually.
    • Step 4: For reactions occurring in acidic medium, add H2O to balance O atoms and H+  to balance H atoms.
    • Step 5: Add electrons to one side of the half reaction to balance the charges. If needed, make the number of electrons equal in the two half reactions by multiplying one or both half reactions by appropriate number.
    • Step 6: Add the two half reactions to achieve the overall reaction and cancel the electrons on each side.
    • Step 7: Verify that the equation contains the same type and number of atoms and the same charges on both sides of the equation.

    Redox Reactions As The Basis For Titration

    In redox systems, the titration method can be adopted to determine the strength of a reductant/oxidant using a redox sensitive indicator.

    Redox Reactions And Electrode Processes: Electrochemical Cells

    • Electrochemical Cell: A device in which the redox reaction is carried indirectly and the decrease in energy appears as the electrical energy are called electrochemical cell.
    • Electrolytic Cell: The cell in which electrical energy is converted into chemical energy. Example, when lead
      storage battery is recharged, it acts as electrolytic cell.
    • Redox Reactions and Electrode Processes: When zinc rod is dipped in copper sulphate solution redox
      reaction begins hence, zinc is oxidised to Zn2+ ions and Cu2+ ions are reduced to metal.
    • Redox Reaction: Reactions in which oxidation and reduction occur simultaneously are called redox
    • Electrode Potential: It is the potential difference between the electrode and its ions in solution.
    • Standard Electrode Potential: It is the potential of an electrode with respect to standard hydrogen electrode.
    • Electrochemical Series: It is activity series. It has been formed by arranging the metals in order of increasing standard reduction potential value.