Is that possible? The problem assigned in class didnt give information about the electrodes, and it asked about the formation of Ag(s). The problem itself was easy, but I realized that this electrolysis can't happen, because of the fact that all the possible reactions are reductions:
NO3(-) + 4H(+) + 3e(-) ----> NO + 2H2O
Ag(+) + e(-) ------> Ag
Besides, E° of NO3(-) is higher, so it should be the one doing the reduction. Maybe it cant do it because there no acid?
Silver nitrate – AgNO3 – melts at 212’C and decomposes at 440’C (Merck Index). So, it would be possible to maintain a liquid vat of the stuff. Electrochemistry does not require a specific chemical reactant for redox reactions. The electrons are provided by the external electrical current. Thus, we only need to consider the potential voltage required to reduce Ag+ to Ag0. Ag+ + e- --> Ag0Eo = 0.80 V
NO3- + 4 H+ + 3 e = NO + 2 H2O . . . Eo = 0.964 V (in water solution)
The lack of any hydrogen to combine with the freed electrons would not be a problem. O2(g) + 4 H+ (aq) + 4 e– 2 H2O(l) Eo = 1.229 This would indicate the conversion TO O2 would contribute -1.229V to the overall system. Whether the NO3 is even reduced may be unnecessary. As long as a current is possible through the liquid and an over-voltage exceeding 0.80V can be applied, it should be possible to plate out solid silver on the cathode. Silver melts at 918’C, so it would not remain in the liquid. An overvoltage of -1.0V might produce silver and nitrogen oxide at the cathode and oxygen at the anode.