Balance the following equations by the oxidation number method.
(i) Fe2+ + H+ + Cr2072- →Cr3+ + Fe3+ + H20
(ii) I2 + N0–3→ N02 +I03
(iii) I2 + S2032- →I– + S4062- ‘
(iv) MnO, + C2042-→ Mn2+ + CO2
Total increase in O.N. = 5×2= 10
Total decrease in O.N. = 1
To equalize O.N. multiply NO3–, by 10
I2 + 10no3–→ 10no2 + IO3–
Balancing atoms other than O and H
I2 + 10no3–→ 10NO2 + 2 IO3–
Balancing O and H
I2 + lO no3– + 8H+→ 10NO2 + 2 IO3– + 4H20
To equalize O.N. multiply C02 by 2. –
Mn02 + C202-4 →Mn2+ + 2CO2
Balance H and O by adding 2H20 on right side, and 4H+ on left side of equation.
MnO2 + C202-4 + 4H+ →Mn2+ + 2CO2 + 2H20
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.
Calculate the oxidation number of phosphorus in the following species.
(a) HPO32- and (b) P043-
Nitric acid is an oxidizing agent and reacts with PbO but it does not react with Pb02. Explain why?
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
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?
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions.
What is meant by electrochemical series? What are characteristics of electrochemical series?
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?
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.
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)
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.
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.
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.
The compound AgF2 is unstable. However, if formed, the compound acts as a very strong oxidising agent. Why?
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 which of the following compounds, an element exhibits two different oxidation states.
(a) NH2OH
(b) NH4NO3
(c) N2H4
(d) N3H
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.
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)?
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
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions.
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.
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?
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
Which of the following elements does not show disproportionation tendency?
(a) Cl
(b) Br
(c) F
(d) I
(a) Balance the following equation by oxidation number method or by ion electron (half reaction) method.
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
Fluorine reacts with ice and results in the change:
H20(S) + F2 (g) ——-> HF(g) + HOF(g)
Justify that this reaction is a redox reaction.
(a)Give two important functions of salt bridge.
(b)Balance the following equation by oxidation number method:
Fe2+ + Cr2O72- + H+ ———> Fe3+ + Cr3+ +H2O
Calculate the oxidation number of each sulphur atom in the following compounds:
(a) Na2S203
(b) Na2S406
(c) Na2S03
(d) Na2S04
Justify-giving reactions that among halogens, fluorine is the best oxidant and among hydrohalic compounds, hydroiodic add is the best reductant.
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 Jour informations about the reaction:
(CN)2(g) + 2OH–(aq) —–> CN–(aq) + CNO–(aq) + H2O(l)
