Redox Reactions

$$2 \mathrm{IO}_{3}^{-}+x \mathrm{I}^{-}+12 \mathrm{H}^{+} \rightarrow 6 \mathrm{I}_{2}+6 \mathrm{H}_{2} \mathrm{O}$$ What is the value of $$x$$ ?

In acidic medium, $\mathrm{K}_2 \mathrm{Cr}_2 \mathrm{O}_7$ shows oxidising action as represented in the half reaction: $$\mathrm{Cr}_2 \mathrm{O}_7{ }^{2-}+\mathrm{XH}^{+}+\mathrm{Ye}^{\ominus} \rightarrow 2 \mathrm{~A}+\mathrm{ZH}_2 \mathrm{O}$$ $\mathrm{X}, \mathrm{Y}, \mathrm{Z}$ and $\mathrm{A}$ are respectively are :

Chlorine undergoes disproportionation in alkaline medium as shown below : $$\mathrm{aCl}_{2(\mathrm{~g})}+\mathrm{b} \mathrm{OH}_{(\mathrm{aq})}^{-} \rightarrow \mathrm{c} \mathrm{ClO}_{(\mathrm{aq)}}^{-}+\mathrm{d} \mathrm{Cl}_{(\mathrm{aq})}^{-}+\mathrm{e} \mathrm{H}_2 \mathrm{O}_{(\mathrm{l})}$$ The values of $$a, b, c$$ and $$d$$ in a balanced redox reaction are respectively :

The volume of 0.1N dibasic acid sufficient to neutralize 1 g of a base that furnishes 0.04 mole of OH^− in aqueous solution is :

The hardness of a water sample (in terms of equivalents of CaCO₃) containing 10^–3 M CaSO₄ is (Molar mass of CaSO₄ = 136 g mol^–1)

In order to oxidise a mixture of one mole of each of FeC₂O₄, Fe₂(C₂O₄)₃, FeSO₄ and Fe₂(SO₄)₃ in acidic medium, the number of moles of KMnO₄ required is :

Which of the following is a redox reaction ?

Which of the following chemical reactions depicts the oxidizing behaviour of H₂SO₄?

Which one of the following is an example of disproportionation reaction ?

In which of the following reactions H₂O₂ acts as a reducing agent? 1. H₂O₂ + 2H⁺ + 2e⁻ $$\to$$ 2H₂O 2. H₂O₂ - 2e⁻ $$\to$$ O₂ + 2H⁺ 3. H₂O₂ + 2e⁻ $$\to$$ 2OH⁻ 4. H₂O₂ + 2OH⁻ - 2e⁻ $$\to$$ O₂ + 2H₂O

In the reaction of oxalate with permanganate in acidic medium, the number of electrons involved in producing one molecule of CO₂ is :

The redox reaction among the following is :

Statement I : Sodium hydride can be used as an oxidising agent. Statement II : The lone pair of electrons on nitrogen in pyridine makes it basic. Choose the CORRECT answer from the options given below :

Which of the following cannot function as an oxidising agent?

In alkaline medium, $$\mathrm{MnO}_4^{-}$$ oxidises $$\mathrm{I}^{-}$$ to

When $$\mathrm{MnO}_2$$ and $$\mathrm{H}_2 \mathrm{SO}_4$$ is added to a salt $$(\mathrm{A})$$, the greenish yellow gas liberated as salt (A) is :

Thiosulphate reacts differently with iodine and bromine in the reactions given below: $$\begin{aligned} & 2 \mathrm{~S}_2 \mathrm{O}_3^{2-}+\mathrm{I}_2 \rightarrow \mathrm{S}_4 \mathrm{O}_6^{2-}+2 \mathrm{I}^{-} \\ & \mathrm{S}_2 \mathrm{O}_3^{2-}+5 \mathrm{Br}_2+5 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{SO}_4^{2-}+4 \mathrm{Br}^{-}+10 \mathrm{H}^{+} \end{aligned}$$ Which of the following statement justifies the above dual behaviour of thiosulphate?

The number of ions from the following that are expected to behave as oxidising agent is : $$\mathrm{Sn}^{4+}, \mathrm{Sn}^{2+}, \mathrm{Pb}^{2+}, \mathrm{Tl}^{3+}, \mathrm{Pb}^{4+}, \mathrm{Tl}^{+}$$

The correct order of the oxidation states of nitrogen in NO, N₂O, NO₂ and N₂O₃ is :