NCERT Solutions: Class 11 Chemistry Chapter 7 - Redox Reactions

In F2, fluorine has oxidation number 0. In HF and HOF, fluorine is -1, so fluorine is reduced. In H2O, oxygen is -2. In HOF, H is +1 and F is -1, so oxygen is 0. Oxygen therefore increases from -2 to 0 and is oxidised. Since oxidation and reduction occur together, the reaction is redox.

Apply oxidation-number sum → charge. In Cr2O7^2-, 2Cr + 7(-2) = -2, so Cr = +6. In NO3-, N + 3(-2) = -1, so N = +5. For H2SO5, the structure is HO-SO2-OOH, containing one peroxide linkage. The two peroxide oxygens are -1 each, while the other three oxygens are -2. Thus 2(+1) + S + 3(-2) + 2(-1) = 0, giving S = +6. If all five oxygens are incorrectly counted as -2, sulphur would appear as +8, which is the fallacy.

Balance cation and anion charges: Hg2+ with two Cl- gives HgCl2; Ni2+ with SO4^2- gives NiSO4; Sn4+ with O2- gives SnO2; two Tl+ ions with SO4^2- give Tl2SO4; two Fe3+ ions with three SO4^2- ions give Fe2(SO4)3; two Cr3+ ions with three O2- ions give Cr2O3.

The examples are chosen so that standard oxidation-number rules give the required sequence. Carbon becomes progressively more positive as H atoms in methane are replaced by electronegative Cl atoms. Nitrogen examples range from hydrides, where N is negative, through elemental N2, to oxides and oxoacids, where N is positive.

Sulphur in SO2 is +4, so it can be oxidised to +6 or reduced to lower states. Oxygen in H2O2 is -1, so it can be oxidised to O2 (0) or reduced to H2O (-2). In O3, oxygen is 0 and readily accepts electrons to form oxide or water species. In HNO3, nitrogen is +5, its maximum common oxidation state, so it cannot be further oxidised and acts as an oxidising agent.

In photosynthesis, the oxygen evolved as O2 comes from water, not directly from CO2, so writing water on both sides helps track the oxygen atoms. In the ozone-hydrogen peroxide reaction, the two O2 molecules are not necessarily equivalent in origin, so writing them separately is more informative. The path can be investigated using isotopic labelling, such as H2^18O or ^18O-labelled peroxide/ozone, followed by analysis of the labelled oxygen in products.

Fluorine is highly electronegative and can stabilise Ag(II) only with difficulty. AgF2 contains Ag2+, while Ag+ is much more stable. Therefore AgF2 has a strong tendency to accept an electron and be reduced to AgF, oxidising other substances in the process.

For carbon, limited oxygen gives CO, where carbon is +2, while excess oxygen gives CO2, where carbon is +4: 2C + O2 → 2CO and C + O2 → CO2. For phosphorus, limited chlorine gives PCl3, where P is +3, while excess chlorine gives PCl5, where P is +5: P4 + 6Cl2 → 4PCl3 and P4 + 10Cl2 → 4PCl5. For sulphur, limited oxygen gives SO2, where S is +4, while stronger/excess oxidation gives SO3, where S is +6: S + O2 → SO2 and 2SO2 + O2 → 2SO3.

A substance oxidised is the reducing agent; a substance reduced is the oxidising agent. In (a), Ag+ in AgBr is reduced to Ag, while hydroquinone C6H6O2 is oxidised to quinone C6H4O2. In (b) and (c), formaldehyde is oxidised to formate, while Ag+ or Cu2+ is reduced. In (d), nitrogen in hydrazine goes from -2 to 0, so N2H4 is oxidised; oxygen in H2O2 goes from -1 to -2, so H2O2 is reduced. In (e), Pb(0) is oxidised to PbSO4 containing Pb2+, while PbO2 containing Pb4+ is reduced to Pb2+.

Iodine is a milder oxidant, so it oxidises thiosulphate only to tetrathionate, S4O6^2-. Bromine is a stronger oxidant and oxidises thiosulphate further to sulphate, SO4^2-. The different oxidising strengths of I2 and Br2 therefore give different oxidation products from the same reductant.

Fluorine can oxidise other halide ions: F2 + 2Cl- → 2F- + Cl2, F2 + 2Br- → 2F- + Br2, and F2 + 2I- → 2F- + I2. The ease of oxidation of halide ions increases down the group, so I- in HI is the strongest reducing halide. For example, 2HI + Cl2 → 2HCl + I2 and 2HI + Br2 → 2HBr + I2.

In XeO6^4-, xenon is +8. In XeO3, xenon is +6, so xenon is reduced. Fluoride ion changes from -1 in F- to 0 in F2, so it is oxidised. Since XeO6^4- oxidises F- to F2, the compound Na4XeO6 containing this ion behaves as a strong oxidising agent.

Both Ag+ and Cu2+ oxidise the stronger reducing agent H3PO2 to H3PO4, being reduced to Ag and Cu respectively. However, benzaldehyde reduces Ag+ to Ag but does not reduce Cu2+ under the stated conditions. Therefore Ag+ has the greater oxidising tendency and is the stronger oxidising agent here.

Each equation is balanced by equating electron gain and loss, then balancing O and H with H2O, H+ or OH- according to the medium. Charge and atom checks confirm: (a) both sides charge -8; (b) both sides charge -1; (c) both sides charge +6; (d) both sides charge 0.

In (CN)2, carbon is in the +3 oxidation state. In CN-, carbon is +2, so one carbon is reduced. In CNO-, carbon is +4, so the other carbon is oxidised. The same reactant is both oxidised and reduced, giving cyanide and cyanate in basic solution; hence it is disproportionation.

One Mn3+ is reduced to Mn2+ and another Mn3+ is oxidised to Mn(IV) in MnO2. Balancing oxygen with water and hydrogen with H+ gives 2Mn3+ + 2H2O → Mn2+ + MnO2 + 4H+. Atoms and total charge are balanced.

Fluorine is the most electronegative element and shows only -1 in compounds. Caesium is an alkali metal and shows +1. Iodine can show -1 as iodide and positive oxidation states in oxo-compounds. Neon is a noble gas and ordinarily shows neither positive nor negative oxidation states.

Sulphur in SO2 is oxidised from +4 to +6 in H2SO4, while chlorine is reduced from 0 in Cl2 to -1 in HCl. Balancing in water gives Cl2 + SO2 + 2H2O → 2HCl + H2SO4.

The reaction is 4NH3 + 5O2 → 4NO + 6H2O. Moles NH3 = 10.00/17.0 = 0.588 mol. Moles O2 = 20.00/32.0 = 0.625 mol. To consume 0.588 mol NH3, O2 required = (5/4)(0.588) = 0.735 mol, so O2 is limiting. Moles NO formed = (4/5)(0.625) = 0.500 mol. Mass NO = 0.500 x 30.0 = 15.0 g.

With silver electrodes in AgNO3, Ag+ is reduced to Ag at the cathode and Ag metal is oxidised to Ag+ at the anode. With platinum electrodes in AgNO3, Ag+ is reduced at the cathode and water is oxidised at the anode to oxygen. In dilute H2SO4 with platinum electrodes, water/H+ gives H2 at the cathode and water gives O2 at the anode. In CuCl2 with platinum electrodes, Cu2+ is reduced to Cu at the cathode and chloride ions are oxidised to Cl2 at the anode.

A more reactive metal displaces a less reactive metal from its salt solution. From the activity/electrochemical series, the decreasing displacement tendency is Mg, Al, Zn, Fe, Cu.

Zinc is oxidised, so the zinc electrode is the anode and is negatively charged in the galvanic cell. Silver ion is reduced at the cathode. The cell notation is Zn(s)|Zn2+(aq)||Ag+(aq)|Ag(s). Electrons flow through the external wire from Zn to Ag, while ions in the solutions and salt bridge complete the internal circuit.