Consider the following reactions :

2 S₂O₃^2–_(aq) + I_2(s) → S₄O₆^2–_(aq) + 2I^–_(aq)

S₂O₃^2–_(aq) + 2Br_2(l) + 5H₂O_(l) → 2SO₄^2–_(aq) + 4Br^–_(aq) + 10H⁺_(aq)

The same reductant, thiosulphate reacts differently with iodine and bromine because:

1. Br₂ is a stronger oxidizing agent than I₂

2. I₂ is a stronger oxidizing agent than Br₂

3. Both '1' and '2'

4. Neither '1' nor '2'

The correct statement(s) about the given reaction is -

\(\mathrm{{X} eO_{6 (aq)}^{4 -} + 2 F^{- 1}_{(aq)} + 6 H^{+}_{(aq )} \rightarrow}~\mathrm{{X} eO_{3 (g)} + {F}_{2 (g)} + 3 H_{2} O_{(l)}}\)${}_{}$

1. ${\mathrm{XeO}}_6^{4-}$ oxidises F^-

2. The oxidation number of F increases from -1  to  zero

3. ${\mathrm{XeO}}_6^{4-}$ is a stronger oxidizing agent that ${\mathrm{F}}^{-}$

4. All of the above.

Consider the following reactions:

Consider the following reactions:

(a) H₃PO_2(aq) + 4AgNO_3(aq) + 2H₂O_(l) → H₃PO_4(aq) + 4Ag_(s) + 4HNO_3(aq)

(b) H₃PO_2(aq) + 2CuSO_4(aq) + 2H₂O_(l) → H₃PO_4(aq) + 2Cu_(s) + H₂SO_4(aq)

$$\begin{gathered} \left(\mathrm{c}\right) {\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CHO}}_{\left(\mathrm{l}\right) }+ 2{{\left[\mathrm{Ag}{\left({\mathrm{NH}}_3\right)}_2\right]}^{+}}_{\left(\mathrm{aq}\right)} + 3{{\mathrm{OH}}^{-}}_{\left(\mathrm{aq}\right)}\to {\mathrm{C}}_6{\mathrm{H}}_5{{\mathrm{COO}}^{-}}_{\left(\mathrm{aq}\right)} + {\mathrm{Ag}}_{\left(\mathrm{s}\right)} + 4{{\mathrm{NH}}_3}_{\left(\mathrm{aq}\right)} +{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{l}\right)} \\ \left(\mathrm{d}\right) {\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CHO}}_{\left(\mathrm{l}\right)} + 2{{\mathrm{Cu}}^{2+}}_{\left(\mathrm{aq}\right)} + 5{{\mathrm{OH}}^{-}}_{\left(\mathrm{aq}\right)}\to \mathrm{No} \mathrm{change} \mathrm{observed} \end{gathered}$$

Draw the inference about the behaviour of Ag⁺ and Cu²⁺ ions from these reactions 

1. Ag⁺ is a stronger oxidizing agent than Cu²⁺

2. Cu²⁺ is a stronger oxidizing agent than Ag⁺

3. Ag⁺ and Cu²⁺ are reducing agents

4.  Ag⁺ is a reducing agent and Cu²⁺ is an oxidising agent

The oxidising agent and reducing agent in the given reaction are

$5{\mathrm{P}}_{4\left(\mathrm{s}\right)}+12{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{l}\right)}+12{{\mathrm{HO}}^{-}}_{\left(\mathrm{aq}\right)}\to$
$8{\mathrm{PH}}_{3\left(\mathrm{g}\right)}+12{{\mathrm{HPO}}_2^{-}}_{\left(\mathrm{aq}\right)}$

1. Oxidising agent = ${\mathrm{P}}_4$; Reducing agent = ${\mathrm{P}}_4$

2. Oxidising agent = ${\mathrm{P}}_4$; Reducing agent = ${\mathrm{H}}_2\mathrm{O}$

3. Oxidising agent = ${\mathrm{H}}_2\mathrm{O}$; Reducing agent = ${\mathrm{P}}_4$

4. None of the above

The correct statement about the given reaction is-

(CN)_2(g) + 2OH^-_(aq) $\mathbf{\to }$CN^-_(aq) + CNO^-_(aq) + H₂O_(l)

The ${\mathrm{Mn}}^{3+}$ ion is unstable in solution and undergoes disproportionation reaction to give ${\mathrm{Mn}}^{2+},$ ${\mathrm{MnO}}_2$ $\mathrm{and}$ ${\mathrm{H}}^{+}$ ion. The balanced ionic equation for the reaction is-

Which element exhibits both positive and negative oxidation states?
1. Cs
2. Ne
3. I
4. F

The balanced equation for the reaction between chlorine and sulphur dioxide in water is:

The set of metals that can show disproportionation reaction is:
1. Cu, Na, Li
2. Mg, F, Ne
3. P, Cl, S
4. Mn, Cu, Ga

The maximum weight of nitric oxide that can be obtained starting only with 10.00 g of ammonia and 20.00 g of oxygen is: