Redox Reactions

In the reactions given below, identify the species undergoing oxidation and reduction: (i) H₂S(g) + Cl₂(g) → 2HCl(g) + S(s) (ii) 3Fe₃O₄(s) + 8Al(s) → 9Fe(s) + 4Al₂O₃(s) (iii) 2Na(s) + H₂(g) → 2NaH(s)

Justify that the reaction: 2Na(s) + H₂(g) → 2NaH(s) is a redox change.

Using Stock notation, represent the following compounds: HAuCl₄, Tl₂O, FeO, Fe₂O₃, CuI, CuO, MnO and MnO₂.

Justify that the reaction: 2Cu₂O(s) + Cu₂S(s) → 6Cu(s) + SO₂(g) is a redox reaction. Identify the species oxidised/reduced, which acts as an oxidant and which acts as a reductant.

Which of the following species, do not show disproportionation reaction and why? ClO⁻, ClO₂⁻, ClO₃⁻ and ClO₄⁻. Also write reaction for each of the species that disproportionates.

Suggest a scheme of classification of the following redox reactions: (a) N₂(g) + O₂(g) → 2 NO(g) (b) 2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g) (c) NaH(s) + H₂O(l) → NaOH(aq) + H₂(g) (d) 2NO₂(g) + 2OH⁻(aq) → NO₂⁻(aq) + NO₃⁻(aq) + H₂O(l)

Why do the following reactions proceed differently? Pb₃O₄ + 8HCl → 3PbCl₂ + Cl₂ + 4H₂O and Pb₃O₄ + 4HNO₃ → 2Pb(NO₃)₂ + PbO₂ + 2H₂O

Write the net ionic equation for the reaction of potassium dichromate(VI), K₂Cr₂O₇ with sodium sulphite, Na₂SO₃, in an acid solution to give chromium(III) ion and the sulphate ion.

Permanganate ion reacts with bromide ion in basic medium to give manganese dioxide and bromate ion. Write the balanced ionic equation for the reaction.

Permanganate(VII) ion, MnO₄⁻ in basic solution oxidises iodide ion, I⁻ to produce molecular iodine (I₂) and manganese (IV) oxide (MnO₂). Write a balanced ionic equation to represent this redox reaction.

Justify that the following reactions are redox reactions: (a) CuO(s) + H₂(g) → Cu(s) + H₂O(g) (b) Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g) (c) 4BCl₃(g) + 3LiAlH₄(s) → 2B₂H₆(g) + 3LiCl(s) + 3AlCl₃(s) (d) 2K(s) + F₂(g) → 2K⁺F⁻(s) (e) 4 NH₃(g) + 5 O₂(g) → 4NO(g) + 6H₂O(g)

Fluorine reacts with ice and results in the change: H₂O(s) + F₂(g) → HF(g) + HOF(g) Justify that this reaction is a redox reaction.

Write formulas for the following compounds: (a) Mercury(II) chloride (b) Nickel(II) sulphate (c) Tin(IV) oxide (d) Thallium(I) sulphate (e) Iron(III) sulphate (f) Chromium(III) oxide

While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why?

Consider the reactions: (a) 6 CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g) (b) O₃(g) + H₂O₂(l) → H₂O(l) + 2O₂(g) Why it is more appropriate to write these reactions as: (a) 6CO₂(g) + 12H₂O(l) → C₆H₁₂O₆(aq) + 6H₂O(l) + 6O₂(g) (b) O₃(g) + H₂O₂(l) → H₂O(l) + O₂(g) + O₂(g) Also suggest a technique to investigate the path of the above (a) and (b) redox reactions.

The compound AgF₂ is unstable compound. However, if formed, the compound acts as a very strong oxidising agent. Why?

Whenever a reaction between an oxidising agent and a reducing agent is carried out, a compound of lower oxidation state is formed if the reducing agent is in excess and a compound of higher oxidation state is formed if the oxidising agent is in excess. Justify this statement giving three illustrations.

Identify the substance oxidised reduced, oxidising agent and reducing agent for each of the following reactions: (a) 2AgBr (s) + C₆H₆O₂(aq) → 2Ag(s) + 2HBr (aq) + C₆H₄O₂(aq) (b) HCHO(l) + 2[Ag (NH₃)₂]⁺(aq) + 3OH⁻(aq) → 2Ag(s) + HCOO⁻(aq) + 4NH₃(aq) + 2H₂O(l) (c) HCHO(l) + 2 Cu²⁺(aq) + 5 OH⁻(aq) → Cu₂O(s) + HCOO⁻(aq) + 3H₂O(l) (d) N₂H₄(l) + 2H₂O₂(l) → N₂(g) + 4H₂O(l) (e) Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

Consider the reactions: 2 S₂O₃²⁻ (aq) + I₂(s) → S₄O₆²⁻ (aq) + 2I⁻(aq) S₂O₃²⁻ (aq) + 2Br₂(l) + 5 H₂O(l) → 2SO₄²⁻ (aq) + 4Br⁻(aq) + 10H⁺(aq) Why does the same reductant, thiosulphate react differently with iodine and bromine?

Justify giving reactions that among halogens, fluorine is the best oxidant and among hydrohalic compounds, hydroiodic acid is the best reductant.