Balance the following oxidation-reduction reactions that occur in acidic solution using the half-reaction method. (Use the lowest possible coefficients. Be sure to specify states such as (aq) or (s). If a box is not needed, leave it blank. Add H₂O or H* to any side of the reaction if it is needed. Fe²+ (aq) + MnO4 (aq) → Fe³+ (aq) + Mn²+ (aq) a. b. - → Fe²+ (aq) + MnO4 (aq) + Fe³+ (aq) + Mn²+ (aq) + As2O3 (8) + NO3(aq) → H3 AsO4 (aq) + NO(g) As₂O3 (8) + NO3(aq) + H3 AsO4 (aq) + + NO(g) + + + +

For the following (potentially unbalanced) reaction, determine what is being (a) reduced and (b) oxidized, (c) write the redox half-reactions, and (d) write a balanced net ionic equation. ZnCl2(aq) + Ti(s) −−→ Zn(s) + TiCl4(aq)

Balance the following half-reactions by adding the appropriate number of electrons (e-). Then, classify each reaction as an oxidation or reduction half-reaction. Note that for each of the four reactions, one of the gray boxes will be left blank and the other will be filled with electron(s). Use the symbol e- to represent an electron. Part 4 + 12(s) + 6H2O(1) > 2I0-3 + 12H+(aq) + --

Balance the following redox reaction if it occurs in basic solution. What are the coefficients in front of Cl2 and OH- in the balanced reaction?Cl2(g) → ClO3-(aq) + Cl-(aq)

Write balanced half-reactions for the following redox reaction: 1₂(s)+2 NO₂(g) + 2 H₂O(1) → 21 (aq) + 2NO3(aq) + 4H* (aq) reduction: oxidation: 0 ロ→ロ X

Write balanced half-reactions for the following redox reaction:

The hydroperoxide ion, HO2–(aq), reacts with permanganate ion, MnO4–(aq) to produce MnO2(s) and oxygen gas. Balance the equation for the oxidation of hydroperoxide ion to O2(g) by permanganate ion in a basic solution.

For the following (potentially unbalanced) reaction, determine what is being (a) reduced and (b) oxidized, (c) write the redox half-reactions, and (d) write a balanced net ionic equation. FeI2(aq) + CuI2(aq) −−→ CuI(s) + FeI3(aq)