Chapter 8: Redox Reactions

Ans: The given reaction is a redox change. 

\[\text{2Na}\left( \text{s} \right)\text{+}{{\text{H}}_{\text{2}}}\left( \text{g} \right)\to \text{2NaH}\left( \text{s} \right)\text{ }\!\!~\!\!\text{ }\]

The half reaction is:

\[\text{2Na}\left( \text{s} \right)\to \text{2N}{{\text{a}}^{\text{+}}}\left( \text{g} \right)\text{+2}{{\text{e}}^{\text{-}}}\]

The other half reaction is:

\[{{\text{H}}_{\text{2}}}\left( \text{g} \right)\text{+2}{{\text{e}}^{\text{-}}}\to \text{2}{{\text{H}}^{\text{-}}}\left( \text{g} \right)\]

This splitting of the reaction into two half-reactions automatically reveals here that sodium is oxidized, and hydrogen is reduced. Any substance which loses electron is oxidized and gains electron is reduced hence is the case of sodium and hydrogen atoms respectively. Hence, the complete reaction is a redox change.

Ans: On the basis of standard reduction potential suggested in the reactivity series (ii) reaction can take place as Mg has more negative value of \[{{\text{E}}^{\odot }}\]cell. Hence, Mg will be oxidized by losing electron and iron will be reduced by gaining electron.

Ans: Disproportionate is defined as the reaction in which one compound of intermediate oxidation state converts to two compounds, one of higher and one of lower oxidation states So, to occur such type of redox reaction, the element should exist in at least three oxidation states. So that element present in the intermediate state and it can change to both higher and lower oxidation state during disproportionate reaction. Fluorine is the most electronegative element and a strong oxidizing agent and is the smallest in size of all the halogens. It does not show a positive oxidation state (shows only −1 oxidation state) and hence, does not undergo disproportionate reaction.

Ans: The given reactions are as given below-

\[\text{ Cu + Z}{{\text{n}}^{\text{2+}}}\to \text{ C}{{\text{u}}^{\text{2+}}}\text{ + Zn}\]

\[\text{ Mg + F}{{\text{e}}^{\text{2+}}}\to \text{ M}{{\text{g}}^{\text{2+}}}\text{ + Fe}\]

\[\text{ B}{{\text{r}}_{\text{2}}}\text{ + 2C}{{\text{l}}^{\text{--}}}\to \text{ C}{{\text{l}}_{\text{2}}}\text{ + 2B}{{\text{r}}^{\text{--}}}\]

\[\text{ Fe + C}{{\text{d}}^{\text{2+}}}\to \text{ Cd + F}{{\text{e}}^{\text{2+}}}\]

Redox couples are given as- \[\text{C}{{\text{u}}^{\text{2+}}}\text{/Cu }\!\!~\!\!\text{ }\] and \[\text{Z}{{\text{n}}^{\text{2+}}}\text{/Zn}\].

 \[\text{M}{{\text{g}}^{\text{2+}}}\text{/Mg}\] and \[\text{F}{{\text{e}}^{\text{2+}}}\text{/Fe}\].

\[\text{B}{{\text{r}}_{\text{2}}}\text{/B}{{\text{r}}^{\text{-}}}\] and \[\text{C}{{\text{l}}_{\text{2}}}\text{/C}{{\text{l}}^{\text{-}}}\].

\[\text{F}{{\text{e}}^{\text{2+}}}\text{/Fe}\] and \[\text{C}{{\text{d}}^{\text{2+}}}\text{/Cd}\].

Ans: We can calculate the oxidation states by - ]

NaClO4 Oxidation no. of chlorine = +7 

Suppose oxidation number of chlorine is x then, 1 + x + 4 × (−2) = 0

∴ x - 7 = 0 

x = +7

We can calculate, the oxidation states, as given below-

NaClO3 Oxidation no. of chlorine = +5

​NaClO Oxidation no. of chlorine = +1

​KClO2 Oxidation no. of chlorine = +3

​Cl2O7 Oxidation no. of chlorine = +7

ClO3 Oxidation no. of chlorine = +6

Cl2O Oxidation no. of chlorine = +1 

NaCl Oxidation no. of chlorine = −1

Cl2 Oxidation no. of chlorine = 0

ClO2 Oxidation no. of chlorine = +4.

Oxidation state (+2) is not present in any of the above compounds.

Ans: We can Measure the electrode potential of the given species by connecting the redox couple of the given species with standard hydrogen electrode. If it is positive, the electrode of the given species acts as reductant and if it is negative, it acts as an oxidant.  Find the electrode potentials of the other given species in the same way, compare the values and determine their comparative strength as an reductant or oxidant.  Example Measurement of standard electrode potential of electrode \[E_{\text{Z}{{\text{n}}^{\text{2+}}}\text{/Zn}}^{\odot }\] using SHE as a reference electrode.

The different types of Redox Reactions are:

• Decomposition Reaction

• Combination Reaction

• Displacement Reaction

• Disproportionation Reactions