Solution of Chapter 8. Redox reactions (Chemistry Part-II Book)

Assign oxidation number to the underlined elements in each of the following species:

NaH₂PO₄

Assign oxidation number to the underlined elements in each of the following species:

NaHSO₄

Assign oxidation number to the underlined elements in each of the following species:

H₄P₂O₇

Assign oxidation number to the underlined elements in each of the following species:

K₂MnO₄

Assign oxidation number to the underlined elements in each of the following species:

CaO₂

Assign oxidation number to the underlined elements in each of the following species:

NaBH₄

Assign oxidation number to the underlined elements in each of the following species:

H₂S₂O₇

Assign oxidation number to the underlined elements in each of the following species:

KAl(SO₄)₂.12 H₂O

What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?

KI₃

What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?

H₂S₄O₆

What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?

Fe₃O₄

What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?

CH₃CH₂OH

What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?

CH₃COOH

Justify that the following reactions are redox reactions:

CuO(s) + H₂(g) → Cu(s) + H₂O(g)

Justify that the following reactions are redox reactions:

Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g)

Justify that the following reactions are redox reactions:

4BCl₃(g) + 3LiAlH₄(s) → 2B₂H₆(g) + 3LiCl(s) + 3 AlCl₃ (s)

Justify that the following reactions are redox reactions:

2K(s) + F₂ (g) → 2K ⁺ F^–(s)

Justify that the following reactions are redox reactions:

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.

Calculate the oxidation number of sulphur, chromium, and nitrogen in H₂SO₅, Cr₂O^2–₇ and NO^–₃. Suggest structure of these compounds. Count for the fallacy.

Write formulas for the following compounds:

Mercury(II) chloride

Write formulas for the following compounds:

Nickel(II) sulphate

Write formulas for the following compounds:

Tin(IV) oxide

Write formulas for the following compounds:

Thallium(I) sulphate

Write formulas for the following compounds:

Iron(III) sulphate

Write formulas for the following compounds:

Chromium(III) oxide

Suggest a list of the substances where carbon can exhibit oxidation states from–4 to + 4 and nitrogen from –3 to + 5.

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) → C₆ H₁₂ O₆(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) → C₆ H₁₂ O₆(aq) + 6H₂O(l) + 6O₂(g)

(b) O₃(g) + H₂O₂(l) → H₂O(l) + 2O₂(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 oxidizing agent. Why?

Whenever a reaction between an oxidizing 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 oxidizing agent is in excess. Justify this statement giving three illustrations.

How do you count for the following observations?

Though alkaline potassium permanganate and acidic potassium permanganate both are used as oxidants, yet in the manufacture of benzoic acid from toluene we use alcoholic potassium permanganate as an oxidant. Why? Write a balanced redox equation for the reaction.

How do you count for the following observations?

When concentrated sulphuric acid is added to an inorganic mixture containing chloride, we get colourless pungent smelling gas HCl, but if the mixture contains bromide then we get red vapour of bromine. Why?

Identify the substance oxidized reduced, oxidizing agent and a reducing agent for each of the following reactions:

2AgBr (s) + C₆H₆O₂ (aq) → 2Ag(s) + 2HBr (aq) + C₆H₄O₂(aq)

Identify the substance oxidized reduced, oxidizing agent and a reducing agent for each of the following reactions:

HCHO(l) + 2[Ag (NH₃)₂] ⁺ (aq) + 3OH^– (aq) → 2Ag(s) + HCOO^– (aq) + 4NH₃(aq) + 2H₂O(l)

Identify the substance oxidized reduced, oxidizing agent and a reducing agent for each of the following reactions:

HCHO (l) + 2 Cu²⁺ (aq) + 5 OH^– (aq) → Cu₂O(s) + HCOO^–(aq) + 3H₂O(l)

Identify the substance oxidized reduced, oxidizing agent and a reducing agent for each of the following reactions:

N₂H₄ (l) + 2H₂O₂ (l) → N₂(g) + 4H₂O(l)

Identify the substance oxidized reduced, oxidizing agent and a reducing agent for each of the following reactions:

Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

Consider the reactions:

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.

Why does the following reaction occur?

What conclusion about the compound Na₄XeO₆ (of which is a part) can be drawn from the reaction.

Consider the reactions:

(a) H₃PO₂(aq) + 4 AgNO₃(aq) + 2 H₂O(l) → H₃PO₄(aq) + 4Ag(s) + 4HNO₃(aq)

(b) H₃PO₂(aq) + 2CuSO₄(aq) + 2 H₂O(l) → H₃PO₄(aq) + 2Cu(s) + H₂SO₄(aq)

(c) C₆H₅CHO(l) + 2[Ag (NH₃)₂]+(aq) + 3OH^– (aq) → C₆H₅COO^–(aq) + 2Ag(s) +4NH₃ (aq) + 2 H₂O(l)

(d) C₆H₅CHO(l) + 2Cu²⁺(aq) + 5OH^–(aq) → No change observed.

What inference do you draw about the behaviour of Ag⁺and Cu²⁺from these reactions?

Balance the following redox reactions by ion – electron method:

(in basic medium)

Balance the following redox reactions by ion – electron method:

(in acidic solution)

Balance the following redox reactions by ion – electron method:

H₂O₂ (aq) + Fe²⁺ (aq) → Fe³⁺ (aq) + H₂O (l) (in acidic solution)

Balance the following redox reactions by ion – electron method:

(in acidic solution)

Balance the following equations in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.

P₄(s) + OH^–(aq) → PH₃(g) + HPO^–₂(aq)

Balance the following equations in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.

N₂H₄(l) + ClO^–₃(aq) → NO(g) + Cl^–(g)

Balance the following equations in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.

Cl₂O₇(g) + H₂O₂(aq) → ClO^–₂(aq) + O₂(g) + H⁺

What sorts of informations can you draw from the following reaction?

(CN)₂(g) + 2OH^–(aq) → CN^–(aq) + CNO^–(aq) + H₂O(l)

The Mn³⁺ ion is unstable in solution and undergoes disproportionation to give Mn²⁺, MnO₂, and H⁺ ion. Write a balanced ionic equation for the reaction.

Consider the elements:

Cs, Ne, I and F

(a) Identify the element that exhibits only negative oxidation state.

(b) Identify the element that exhibits only positive oxidation state.

(c) Identify the element that exhibits both positive and negative oxidation states.

(d) Identify the element which exhibits neither the negative nor does the positive oxidation state.

Chlorine is used to purify drinking water. Excess of chlorine is harmful. The excess of chlorine is removed by treating with Sulphur dioxide. Present a balanced equation for this redox change taking place in water.

Refer to the periodic table given in your book and now answer the following questions:

(a) Select the possible non metals that can show disproportionation reaction.

(b) Select three metals that can show disproportionation reaction.

In Ostwald’s process for the manufacture of nitric acid, the first step involves the oxidation of ammonia gas by oxygen gas to give nitric oxide gas and steam.

What is the maximum weight of nitric oxide that can be obtained starting only with 10.00 g. of ammonia and 20.00 g of oxygen?

Using the standard electrode potentials given in the Table 8.1, predict if the reaction between the following is feasible:

Fe³⁺(aq) and I^–(aq)

Using the standard electrode potentials given in the Table 8.1, predict if the reaction between the following is feasible:

Ag⁺(aq) and Cu(s)

Using the standard electrode potentials given in the Table 8.1, predict if the reaction between the following is feasible:

Fe³⁺(aq) and Cu(s)

Using the standard electrode potentials given in the Table 8.1, predict if the reaction between the following is feasible:

Ag(s) and Fe³⁺(aq)

Using the standard electrode potentials given in the Table 8.1, predict if the reaction between the following is feasible:

Br₂(aq) and Fe²⁺(aq).

Predict the products of electrolysis in each of the following:

(i) An aqueous solution of AgNO₃ with silver electrodes

(ii) An aqueous solution AgNO₃ with platinum electrodes

(iii) A dilute solution of H₂SO₄ with platinum electrodes

(iv) An aqueous solution of CuCl₂ with platinum electrodes.

Arrange the following metals in the order in which they displace each other from the solution of their salts.

Al, Cu, Fe, Mg and Zn.

Given the standard electrode potentials,

K⁺/K = –2.93V, Ag⁺/Ag = 0.80V,

Hg²⁺/Hg = 0.79V

Mg²⁺/Mg = –2.37V. Cr³⁺/Cr = –0.74V

arrange these metals in their increasing order of reducing power.

Depict the galvanic cell in which the reaction Zn(s) + 2Ag⁺(aq) → Zn²⁺(aq) +2Ag(s) takes place, Further show:

(i) which of the electrode is negatively charged,

(ii) the carriers of the current in the cell, and

(iii) individual reaction at each electrode.