Hydrogen peroxide, H₂O₂, is a nontoxic bleaching agent being used as a replacement for chlorine in industry and home laundries. The bleaching process is an oxidation, and when hydrogen peroxide acts as a bleaching agent, the only waste it generates is H₂O.

(a) Draw the Lewis structure of hydrogen peroxide and determine the formal charge on each atom. What is the oxidation number of oxygen in hydrogen peroxide? Which is more useful in predicting the ability of H₂O₂ to act as an oxidizing agent, formal charge or oxidation number? Explain your reasoning.

(b) Predict the bond angles at each O atom in H₂O₂. Are all the atoms in the same plane? Is the molecule polar or nonpolar? Explain your reasoning.

(c) Write the valence electron configuration of (i) O₂; (ii) O₂^–; (iii) O₂⁺ ; (iv) O₂^2–. For each species, give the expected bond order and indicate which, if any, are paramagnetic.

(d) The following bond lengths have been reported: (i) O₂, 121 pm; (ii) O₂^–, 134 pm; (iii) O₂⁺, 112 pm; (iv) O₂^2–, 149 pm. Suggest a reason for the differences based on the configurations in part (c).

(e) One reaction in which H₂O₂ acts as an oxidizing agent is Fe²⁺(aq) + H₂O₂ (aq) + H⁺ (aq) → Fe³⁺ (aq) + H₂O(l). Balance the equation and determine what mass of iron(II) can be oxidized to iron(III) by 43.2 mL of 0.200 m H₂O₂(aq).

(f) Hydrogen peroxide can also act as a reducing agent, as in Fe³⁺ (aq) + H₂O₂ (aq) + OH^– (aq) → Fe²⁺ (aq) + H₂O(l) + O₂ (g). Balance the equation and determine what mass of iron(III) can be reduced to iron(II) by 41.8 mL of 0.200 m H₂O₂ (aq).

(g) Hydrogen peroxide must be kept in brown bottles because, in the presence of light, it can disproportionate, which means that it oxidizes and reduces itself in the reaction 2 H₂ O₂ (aq) → 2 H₂ O(l) + O₂(g). How many electrons are transferred in the reaction represented by this equation?

(h) Although peroxides do not generate hazardous waste, they can cause problems in the atmosphere. For example, if a hydrogen peroxide molecule makes its way to the stratosphere, it can break into two · OH radicals, which threaten the ozone layer that protects Earth from harmful radiation. Use data in Table 2D.2 to calculate the minimum frequency and corresponding wavelength of light needed to break the HO—OH bond.

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TABLE 2D.2 Average Bond Dissociation Energies Average bond dissociation energy/(kJ mol) Bond C-H C-C C=C CC* C=C C-O C=O C-N C-F C-Cl C-Br 412 348 612 518 837 360 743 305 484 338 276 Bond C-I N-H N-N N=N N-O N=O N-F N-Cl O-H 0-0 Average bond dissociation energy/(kJ. mol) 238 388 163 409 210 630 270 200 463 157