Which of the following is not an example of redox reaction?

The more positive the value of $E^{\ominus}$, the greater is the tendency of the species to get reduced. Using the standard electrode potential of redox couples given below find out which of the following is the strongest oxidising agent.

$\begin{aligned} & E^{\ominus} \text { values: } \mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}=+0.77 \\ & \qquad \mathrm{I}_2(s) / \mathrm{I}^{-}=+0.54 ; \\ & \mathrm{Cu}^{2+} / \mathrm{Cu}=+0.34 ; \mathrm{Ag}^{+} / \mathrm{Ag}=0.80 \mathrm{~V}\end{aligned}$

$E^{\ominus}$ values of some redox couples are given below. On the basis of these values choose the correct option.

$$\begin{aligned} E^{\ominus} \text { values: } \mathrm{Br}_2 / \mathrm{Br}^{-} & =+1.90 \\ \mathrm{Ag}^{+} / \mathrm{Ag}(s) & =+0.80 \\ \mathrm{Cu}^{2+} / \mathrm{Cu}(s) & =+0.34 ; \mathrm{I}_2(s) / \mathrm{I}^{-}=+0.54 \end{aligned}$$

Using the standard electrode potential, find out the pair between which redox reaction is not feasible.

$$\begin{aligned} & E^{\ominus} \text { values: } \mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}=+0.77 ; \mathrm{I}_2 / \mathrm{I}^{-}=+0.54 ; \\ & \mathrm{Cu}^{2+} / \mathrm{Cu}=+0.34 ; \mathrm{Ag}^{+} / \mathrm{Ag}=+0.80 \mathrm{~V} \end{aligned}$$

Thiosulphate reacts differently with iodine and bromine in the reactions given below

$$\begin{gathered} 2 \mathrm{~S}_2 \mathrm{O}_3^{2-}+\mathrm{I}_2 \rightarrow \mathrm{S}_4 \mathrm{O}_6^{2-}+2 \mathrm{I}^{-} \\ \mathrm{S}_2 \mathrm{O}_3^{2-}+2 \mathrm{Br}_2+5 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{SO}_4^{2-}+2 \mathrm{Br}^{-}+10 \mathrm{H}^{+} \end{gathered}$$

Which of the following statements justifies the above dual behaviour of thiosulphate?

The oxidation number of an element in a compound is evaluated on the basis of certain rules. Which of the following is incorrect in this respect?

In which of the following compounds, an element exhibits two different oxidation states?

Which of the following arrangements represent increasing oxidation number of the central atom?

The largest oxidation number exhibited by an element depends on its outer electronic configuration. With which of the following outer electronic configurations the element will exhibit largest oxidation number?

Identify disproportionation reaction

Which of the following elements does not show disproportionation tendency ?

Assertion (A) Among halogens fluorine is the best oxidant.

Reason (R) Fluorine is the most electronegative atom.

Assertion (A) In the reaction between potassium permanganate and potassium iodide, permanganate ions act as oxidising agent.

Reason (R) Oxidation state of manganese changes from +2 to +7 during the reaction.

Assertion (A) The decomposition of hydrogen peroxide to form water and oxygen is an example of disproportionation reaction.

Reason (R) The oxygen of peroxide is in -1 oxidation state and it is converted to zero oxidation state in $\mathrm{O}_2$ and -2 oxidation state in $\mathrm{H}_2 \mathrm{O}$.

Assertion (A) Redox couple is the combination of oxidised and reduced form of a substance involved in an oxidation or reduction half cell.

Reason (R) In the representation $E_{\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}}^{\ominus}$ and $E_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\ominus}, \mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}$ and $\mathrm{Cu}^{2+} / \mathrm{Cu}$ are redox couples.

Which of the following statement(s) is/are not true about the following decomposition reaction?

$$2 \mathrm{KClO}_3 \longrightarrow 2 \mathrm{KCl}+3 \mathrm{O}_2$$

Identify the correct statement (s) in relation to the following reaction.

$$\mathrm{Zn}+2 \mathrm{HCl} \longrightarrow \mathrm{ZnCl}_2+\mathrm{H}_2$$

The exhibition of various oxidation states by an element is also related to the outer orbital electronic configuration of its atom. Atom (s) having which of the following outermost electronic configurations will exhibit more than one oxidation state in its compounds.

Identify the correct statements with reference to the given reaction

$$\mathrm{P}_4+3 \mathrm{OH}^{-}+3 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{PH}_3+3 \mathrm{H}_2 \mathrm{PO}_2^{-}$$

Which of the following electrodes will act as anodes, which connected to Standard Hydrogen Electrode ?

The reaction $\mathrm{Cl}_2(g)+2 \mathrm{OH}^{-}(a q) \rightarrow \mathrm{ClO}^{-}(a q)+\mathrm{Cl}^{-}(a q)+\mathrm{H}_2 \mathrm{O}(l)$ represents the process of bleaching. Identify and name the species that bleaches the substances due to its oxidising action.

$\mathrm{MnO}_4^{2-}$ undergoes disproportionation reaction in acidic medium but $\mathrm{MnO}_4^{-}$does not. Give reason.

PbO and $\mathrm{PbO}_2$ react with HCl according to following chemical equations

$$\begin{aligned} & 2 \mathrm{PbO}+4 \mathrm{HCl} \longrightarrow 2 \mathrm{PbCl}_2+2 \mathrm{H}_2 \mathrm{O} \\ & \mathrm{PbO}_2+4 \mathrm{HCl} \longrightarrow \mathrm{PbCl}_2+\mathrm{Cl}_2+2 \mathrm{H}_2 \mathrm{O} \end{aligned}$$

Why do these compounds differ in their reactivity ?

Nitric acid is an oxidising agent and reacts with PbO but it does not react with $\mathrm{PbO}_2$. Explain why ?

Write balanced chemical equation for the following reactions.

(a) Permanganate ion $\left(\mathrm{MnO}_4^{-}\right)$reacts with sulphur dioxide gas in acidic medium to produce $\mathrm{Mn}^{2+}$ and hydrogen sulphate ion. (Balance by ion electron method)

(b) Reaction of liquid hydrazine $\left(\mathrm{N}_2 \mathrm{H}_4\right)$ with chlorate ion $\left(\mathrm{ClO}_3^{-}\right)$in basic medium produces nitric oxide gas and chloride ion in gaseous state. (Balance by oxidation number method)

(c) Dichlorine heptaoxide $\left(\mathrm{Cl}_2 \mathrm{O}_7\right)$ in gaseous state combines with an aqueous solution of hydrogen peroxide in acidic medium to give chlorite ion $\left(\mathrm{ClO}_2^{-}\right)$and oxygen gas. (Balance by ion electron method)

Calculate the oxidation number of phosphorus in the following species.

(a) $\mathrm{HPO}_3^{2-}$

(b) $\mathrm{PO}_4^{3-}$

Calculate the oxidation number of each sulphur atom in the following compounds.

(a) $\mathrm{Na}_2 \mathrm{~S}_2 \mathrm{O}_3$

(b) $\mathrm{Na}_2 \mathrm{~S}_4 \mathrm{O}_6$

(c) $\mathrm{Na}_2 \mathrm{SO}_3$

(d) $\mathrm{Na}_2 \mathrm{SO}_4$

Balance the following equations by the oxidation number method.

(a) $\mathrm{Fe}^{2+}+\mathrm{H}^{+}+\mathrm{Cr}_2 \mathrm{O}_7^{-2} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{Fe}^{3+}+\mathrm{H}_2 \mathrm{O}$

(b) $\mathrm{I}_2+\mathrm{NO}_3^{-} \longrightarrow \mathrm{NO}_2+\mathrm{IO}_3^{-}$

(c) $\mathrm{I}_2+\mathrm{S}_2 \mathrm{O}_3^{2-} \longrightarrow \mathrm{I}^{-}+\mathrm{S}_4 \mathrm{O}_6^{2-}$

(d) $\mathrm{MnO}_2+\mathrm{C}_2 \mathrm{O}_4^{2-} \rightarrow \mathrm{Mn}^{2+}+\mathrm{CO}_2$

Identify the redox reaction out of the following reactions and identify the oxidising and reducing agents in them.

(a) $3 \mathrm{HCl}(a q)+\mathrm{HNO}_3(a q) \longrightarrow \mathrm{Cl}_2(g)+\mathrm{NOCl}(g)+2 \mathrm{H}_2 \mathrm{O}(l)$

(b) $\mathrm{HgCl}_2(a q)+2 \mathrm{KI}(a q) \longrightarrow \mathrm{HgI}_2(s)+2 \mathrm{KCl}(a q)$

(c) $\mathrm{Fe}_2 \mathrm{O}_3(\mathrm{~s})+3 \mathrm{CO}(\mathrm{g}) \xrightarrow{\Delta} 2 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_2(\mathrm{~g})$

(d) $\mathrm{PCl}_3(l)+3 \mathrm{H}_2 \mathrm{O}(l) \longrightarrow 3 \mathrm{HCl}(a q)+\mathrm{H}_2 \mathrm{PO}_3(a q)$

(e) $4 \mathrm{NH}_3(a q)+3 \mathrm{O}_2(g) \longrightarrow 2 \mathrm{~N}_2(g)+6 \mathrm{H}_2 \mathrm{O}(\mathrm{g})$

Balance the following ionic equations.

(a) $\mathrm{Cr}_2 \mathrm{O}_7^{2-}+\mathrm{H}^{+}+\mathrm{I}^{-} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{I}_2+\mathrm{H}_2 \mathrm{O}$

(b) $\mathrm{Cr}_2 \mathrm{O}_7^{2-}+\mathrm{Fe}^{2+}+\mathrm{H}^{+} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{Fe}^{3+}+\mathrm{H}_2 \mathrm{O}$

(c) $\mathrm{MnO}_4^{-}+\mathrm{SO}_3^{2-}+\mathrm{H}^{+} \longrightarrow \mathrm{Mn}^{2+}+\mathrm{SO}_4^{2-}+\mathrm{H}_2 \mathrm{O}$

(d) $\mathrm{MnO}_4^{-}+\mathrm{H}^{+}+\mathrm{Br}^{-} \longrightarrow \mathrm{Mn}^{2+}+\mathrm{Br}_2+\mathrm{H}_2 \mathrm{O}$

Match Column I with Column II for the oxidation states of the central atoms.

Match the items in Column I with relevant items in Column II.

Explain redox reaction on the basis of electron transfer. Given suitable examples.

On the basis of standard electrode potential values, suggest which of the following reactions would take place? (Consult the book for $E^{\ominus}$ value)

(a) $\mathrm{Cu}+\mathrm{Zn}^{2+} \longrightarrow \mathrm{Cu}^{2+}+\mathrm{Zn}$

(b) $\mathrm{Mg}+\mathrm{Fe}^{2+} \longrightarrow \mathrm{Mg}^{2+}+\mathrm{Fe}$

(c) $\mathrm{Br}_2+2 \mathrm{Cl}^{-} \longrightarrow \mathrm{Cl}_2+2 \mathrm{Br}^{-}$

(d) $\mathrm{Fe}+\mathrm{Cd}^{2+} \longrightarrow \mathrm{Cd}+\mathrm{Fe}^{2+}$

Why does fluorine not show disproportionation reaction?

Write redox couples involved in the reactions (a) to (d) given in question 34 .

Find out the oxidation number of chlorine in the following compounds and arrange them in increasing order of oxidation number of chlorine. $\mathrm{NaClO}_4, \mathrm{NaClO}_3, \mathrm{NaClO}, \mathrm{KClO}_2, \mathrm{Cl}_2 \mathrm{O}_7, \mathrm{ClO}_3, \mathrm{Cl}_2 \mathrm{O}, \mathrm{NaCl}, \mathrm{Cl}_2, \mathrm{ClO}_2$. Which oxidation state is not present in any of the above compounds?

Which method can be used to find out strength of reductant/oxidant in a solution? Explain with an example.