Question 3 Electrolytic cells require the presence of a power supply to drive the reaction. In most cases, this is done to allow non-spontancous reactions, with negative potentials to take place. In this experiment, we are using a power supply to speed up a spontaneous reaction. The two half-reactions are m(s)âMP"(aq)+xe 2H-(aq) + 2e â H2 (g) oxidation reduction Your unknown metal will be one in the Table, below. In this experiment your unknown is attached to the anode of a power supply, This will drive the oxidation of the metal. The metal is placed in an acidic solution, providing the protons for the reduction reaction. The H2 gas is produced at a metal surface (Cu in the lab experiment) connected to the cathode of the power supply. The gas is generated inside an inverted buret, so that the change in volume of hydrogen gas can be measured during the experiment. You will weigh your unknown metal before and after the experiment, to quantify how many grams of metal have been oxidized. You will measure how many mL of H2 gas are produced and, using the Ideal Gas Law convert this to moles of H2 produced. From this you can calculate the number of moles of electrons exchanged during the electrolysis process. Knowing the grams of metal oxidized, and the number of moles of electrons transferred, you can calculate the equivalent mass of your unknown and compare this to the values in the Table. half-reaction 5° (v) equivalent mass (g/mol e-) Ag + e-Ag +0.80 Fe3+ 3e" Fe-0.036 107.9 18.62 103.6 8.99 12.16 Al3+ + 3e" â Al -1.66 Mg2+ + 2e-Mg| -2.37 Initial mass of metal Click "Weigh Initial Mass of Metal" to view the initial mass of the metal on the balance 21.387
