|
|
|
|
|
"Secondary cells are galvanic cells that can be recharged. Operating them as electrochemical cells such that the electrochemical reactions are reversed recharges them. Primary cells are galvanic cells that cannot be recharged by operating them with a reversed polarity". “I still do not understand this in more than just a ‘qualitative’ way”
If a battery is to be cycled, it is essential that the chemical reactions that take place at the electrodes are efficiently reversed. If only 1% of active material is lost in one cycle, that signifies that 50% capacity would be lost after only fifty cycles. Many, if not most, electrode reactions involve a reconstructive phase change in the crystal chemistry of the active materials. Many possible processes or side reactions leading to battery deterioration or its failure can occur. (more on degradation mode)
“In a lead/acid battery why does the PbSO4 deposit on the electrode instead of just precipitating out as flakes?”
The solubility of lead ions in sulfuric acid is extremely low, preventing the free precipitation of lead sulfate. Moreover the transformation of lead and lead dioxide into lead sulfate at respectively the negative and positive electrodes requires the transfer of electrons which can only occur on the electrodic material itself. The main problem with lead sulfate is its transformation from a useful and reducible sulfate into a much less accessible sulfate in a process called sulfation (read more about this)
“I still don’t really understand why most batteries are considered fully discharged when only 80% of the theoretical capacity has been used. What is the difference (chemically) between ‘regular discharge’ and ‘deep’ discharge?”
“What are Ragone plots and what do they really mean?”
“How is the standard potential for a cell determined that has a ‘multi-step’ process that occurs at one of the electrodes? For example a Zinc/Silver Oxide cell:
|
Anode |
Zn(s) + 2OH- <--> Zn(OH)2 + 2e- |
E0 = -1.25 |
|
Cathode 1 |
2AgO(s) + H2O(l) + 2e- <--> Ag2O(s) + 2OH- |
E0 = 0.57 |
|
Cathode 2 |
Ag2O(s) + H2O(l) + 2e- <--> 2Ag(s) + 2OH- |
E0 = 0.35 |
|
Overall |
AgO(s) + Zn(s) H2O(l) <--> Zn(OH)2 + Ag(s) |
“OCV = ?” |
This question really focuses on the rechargeable Silver/Zinc battery since the primary system is based on the cathode 2 reaction (argentous oxide) and has a flat discharge profile essential for many of its applications.
Well, it really depends where you are in the discharge processes and how fast you have been discharging the battery. In this particular case the first portion of the discharge process will be predominately influenced by cathode 1 reaction (reduction of argentic oxide). Once the higher oxidation form of silver is all reduced to its argentous form the cathode 2 process will control the potential. Usually you reflect on these problems by looking where you are on the polarization curve describing the battery discharge process.
“The standard reduction potentials for a lead/acid, NiCad, and NiMH are 1.69, 1.26, and 1.35 V respectively. If the standard reduction for the ‘gassing’ reaction is lower than all three of them, at 1.23V. Why then does gassing not occur, throughout the charging cycle instead of just when the reactants are depleted? I would predict, that any voltage above 1.23V applied across the terminals would induce gassing, thus causing the re-charging and gassing to happen in parallel.”
