PbO and Pb02 react with HC1 according to following chemical equations:
2PbO + 4HCl → 2PbCl2 + 2H20
Pb02 + 4HC1 → PbCl2 + Cl2 + 2H20
Why do these compounds differ in their reactivity?
In reaction (i), O.N. of none of the atoms undergo a change. Therefore, it is not a redox reaction. It is an acid-base reaction, because PbO is a basic oxide which reacts with HCl acid.
The reaction (ii) is a redox reaction in which Pb02 gets reduced and acts as an oxidizing agent.
Identify the correct statements with reference to the given reaction.
P4 + 30H– + 3H20→ PH3 + 3H2 P0–2
(a) Phosphorus is undergoing reduction only.
(b) Phosphorus is undergoing oxidation only.
(c) Phosphorus is undergoing oxidation as well as reduction.
(d) Hydrogen is undergoing neither oxidation nor reduction
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.
While sulphur dioxide and hydrogen peroxide can act as an oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why?
Predict the products of electrolysis in each of the folloxving:
(i) An aqueous solution of AgNO3 with silver electrodes.
(ii) An aqueous solution of silver nitrate with platinum electrodes.
(iii) A dilute solution of H2S04with platinum electrodes.
(iv) An aqueous solution of CuCl2 with platinum electrodes.
Balance the following redox reactions by ion-electron method.
(a) MnO4–(aq) +I–(aq) ———>Mn02(s) + I2 (s) (in basic medium)
(b) MnO4–(aq) + S02(g) ——-> Mn2+(aq) +H2S04–(in acidic solution)
(c) H2O2(aq) + Fe2+(aq) ———-> Fe3+(aq) + H2O(l) (in acidic solution)
(d) Cr2O72- (aq) + S02 (g)——> Cr3+ (aq) + SO42-(aq) (in acidic solution)
Calculate the oxidation number of phosphorus in the following species.
(a) HPO32- and (b) P043-
Which of the following electrodes will act as anodes, when connected to Standard Hydrogen Electrode?
(a) A13-/A1; E °= -1.66 V
(b) Fe2+ /Fe; E °= -0.44 V
(c) Cu2+/ Cu E °=34 V
(d) F2(g)/2F–(aq) E °= 2.87 V
What is standard hydrogen electrode? For what purpose it is used? What are signs of oxidation potential and reduction potential decided by using SHE (Standard hydrogen electrode)?
Which of the following statement(s) is/are not true about the following decomposition reaction?
2KCIO3 →2KC1 + 302
(a) Potassium is undergoing oxidation.
(b) Chlorine is undergoing oxidation.
(c) Oxygen is reduced.
(d) None of the species are undergoing oxidation or reduction.
Identify the oxidant and the reductant in the following reaction.
N2H4(g) + ClO4(aq) ———–> NO(g) + Cr(aq)
In Ostwald’s process for the manufacture of nitric add, the first step involves the oxidation of ammonia gas by oxygen gas to give nitric oxide gas and steam. What is the maximum wight of nitric oxide that can be obtained starting only with 10.0 g of ammonia and 20.0 g of oxygen?
Nitric acid is an oxidizing agent and reacts with PbO but it does not react with Pb02. Explain why?
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/their compounds.
(a) 3s1
(b) 3dl4s2
(c) 3d24s2
(d) 3s23p3
Which of the following elements does not show disproportionation tendency?
(a) Cl
(b) Br
(c) F
(d) I
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions.
E ° values of some redox complexes are given below. On the basis of these values choose the correct option.
E ° values: Br2/Br– = +1.90; Ag+/Ag(s) = +0.80 Cu2+/Cu(s) = +0.34; I2(s)/I– = +0.54 V
(a) Cu will reduce Br–
(b) Cu will reduce Ag
(c) Cu will reduce I–
(d) Cu will reduce Br2
Justify that the following reactions are redox reactions:
(a) CuO(s) + H2(g) —–> Cu(s) + H20(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)
Consider the reactions:
(a) 6CO2(g) 6H2O(l) ———> C6H12O6(s) + 6O6(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(s) + 6H2O(l) + 6O2(g)
(b) O3(g) + H2O2 (l) ———–> H2O(l) + O2(g) + O2(g)
Also suggest a technique to investigate the path of above (a) and (b) redox reactions.
Identify the redox reactions out of the following reactions and identify the oxidizing and reducing agents in them
Consider the reactions:
(a) H3P02(aq) + 4AgNO3(aq) + 2H2O(l) ————->H3PO4(aq) + 4Ag(s) + 4HNO3(aq)
(b) H3P02(aq) + 2CuS04(aq) + 2H2O(l) ————->H3P04(aq) + 2Cu(s) + H2S04(aq)
(c) C6H5CHO(l) + 2[Ag(NH3)2]+(aq) + 30H–(aq)———–> C6H5COO–(aq) + 2Ag(s) + 4NH3(aq) + 2H20(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?
Write the O.N of all the atoms for the following well known oxidants?
(i) KMnO4 (ii) K2Cr2O7 (iii) KClO4
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions.
Thiosulphate reacts differently with iodine and bromine in the reactions given below:
2S2032_ + I2→S4062- + 2I–
S2032- + 2Br2 + 5H20 →2S042- + 4Br– + 10H+
Which of the following statements justifies the above dual behaviour of thiosulphate?
(a) Bromine is a stronger oxidant than iodine.
(b) Bromine is a weaker oxidant than iodine.
(c) Thiosulphate undergoes oxidation by bromine and reduction by iodine in these reactions.
(d) Bromine undergoes oxidation and iodine undergoes reduction in these reactions.
The compound AgF2 is unstable. However, if formed, the compound acts as a very strong oxidising agent. Why?
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 current in the cell and
(iii) individual reaction at each electrode.
Identify disproportionation reaction
(a) CH4 + 202 → C02 + 2H20
(b) CH4 + 4C12 → CC14 + 4HCl
(c) 2F2 + 20H–→2F– + OF2 + H20
(d) 2N02 + 20H– → N02 + NO–3 + H20
The reaction Cl2(g) + 20H–(aq)→ Cl0–(aq) + Cl–(aq) + H20(l) represents the process of bleaching. Identify and name the species that bleaches the substances due to its oxidizing action.
Consider the elements: Cs, Ne, I, F
(a) Identify the element that exhibits -ve oxidation state.
(b) Identify the element that exhibits +ve oxidation state.
(c) Identify the element that exhibits both +ve and -ve oxidation states.
(d) Identify the element which neither exhibits -ve nor +ve oxidation state.
Consider a voltaic cell constructed with the following substances:
(a) Which substances are oxidised and reduced in this cell?
(b) Which are the negative and positive electrode?
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.
Given the standard electrode potentials,
K+/K = -2.93 V, Ag+/Ag = 0.80 V, Hg2+/Hg = 0.79 V, Mg2+/Mg = -2.37 V,
Cr3+/Cr = -0.74 V. Arrange these metals in increasing order of their reducing power.
(a) Balance the following equation by oxidation number method or by ion electron (half reaction) method.
Identify the correct statement(s) in relation to the following reaction:
Zn + 2HCl → ZnCl2 + H2
(a) Zinc is acting as an oxidant.
(b) Chlorine is acting as a reductant.
(c) Hydrogen ion is acting as an oxidant.
(d) Zinc is acting as a reductant.
PbO and Pb02 react with HC1 according to following chemical equations:
2PbO + 4HCl → 2PbCl2 + 2H20
Pb02 + 4HC1 → PbCl2 + Cl2 + 2H20
Why do these compounds differ in their reactivity?
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 show disproportionation reaction.
