Electron Transfer: Redox Titration and Equivalence Points Quiz

While acid-base titrations involve the transfer of protons (H+), redox titrations are defined by the transfer of electrons between the oxidizing agent and the reducing agent. This change in oxidation states is the driving force of the reaction monitored during the process.

Potassium Permanganate is an intense purple color. During the titration, as it is reduced to Mn2+, it becomes colorless. The first permanent pale pink tinge in the flask signifies that all the analyte has reacted, marking the end point without needing an added chemical indicator.

This is true. In titrations involving iodine, starch is added near the end point. It reacts with the remaining I3- to form a dark blue-black complex. When the blue color disappears, it indicates that all the iodine has been reduced to iodide (I-).

This is the standard reduction half-reaction used in acidic redox titrations. It shows the gain of 5 electrons and the consumption of 8 protons to produce the colorless manganese(II) ion and water.

The oxidizing agent facilitates the oxidation of the analyte by accepting electrons. Consequently, the oxidizing agent itself undergoes reduction, resulting in a decrease in its oxidation state.

In MnO4-, oxygen is assigned -2. Since the overall charge is -1, the calculation is x + 4(-2) = -1, which gives x = +7. During the titration in acidic medium, it is typically reduced to Mn2+ (+2).

The reduction of MnO4- to Mn2+ requires a large supply of protons (8 H+ per MnO4-). Without an acidic environment, the reaction might produce MnO2 (a brown precipitate), which interferes with the visual end point.

Sodium thiosulfate is a standard reducing agent. It reacts with iodine to form tetrathionate ions (S4O6 2-) and iodide ions (I-). This reaction is widely used to determine the concentration of various oxidizing agents indirectly.

This is true. Potentiometric titrations measure the change in electrode potential as the titrant is added. At the equivalence point, there is a sharp change in potential, which can be graphed to determine the volume of titrant used very accurately.

Moles of MnO4- = 0.02 x 0.020 = 0.0004. Moles of Fe2+ = 0.0004 x 5 = 0.002. Concentration = 0.002 / 0.025 L = 0.080 M. The 1:5 stoichiometry is vital for this calculation.

This is false. The equivalence point is the theoretical point where the reactants are in exact stoichiometric proportions. The end point is the physical point where the indicator changes color. A well-designed titration ensures these two points are as close as possible.

This is false. Unlike simple 1:1 acid-base reactions, redox stoichiometry depends on the number of electrons transferred. For example, 1 mole of MnO4- (5 electrons) reacts with 5 moles of Fe2+ (1 electron each). The equivalence point is reached when the moles of electrons lost equals the moles of electrons gained.

For a successful titration, the reaction needs to occur almost instantaneously upon mixing, follow a known stoichiometry, and provide a visible or measurable signal when the equivalence point is reached.

Potassium dichromate contains Cr(VI) which is orange. When it is reduced to Cr(III) during a redox reaction, the solution turns green. This distinct color shift can be used to monitor the progress of the titration.

Potassium dichromate (K2Cr2O7) and Iodine (I2) are classic oxidizing agents. Sodium thiosulfate and oxalic acid are typically used as reducing agents in these analytical procedures.