> -----Original Message----- > From: [email protected] [mailto:[email protected]] > Sent: Friday, 22 June 2001 10:13 > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx > Ivan: Well, then how did Ag+ (by far THE most stable ion of > silver) get to > Ag++, and what is present to allow it to remain in the +2 > oxidation state? > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx
Roger, By giving an electron to two clorine atoms i guess. I don't know how stable this styff is, but I should think that one will find a certain concentration at any time, but like water molecules, the Ag++ is constantly devolving and reforming. > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx > Ivan: OK. Let's look at a scenario that takes into account the > formation of > molecular AgCl(aq) after AgCl(s) has already formed. Add AgCl to > water until > it's saturated. Now add NaCl solution. Due to the common ion > effect, the Ag+ > conc. will continue to decrease until Cl- reaches a concentration > sufficient > to produce AgCl(aq) which represents molecular NOT crystalline > AgCl. Now how > can the Ag+ concentration increase when AgCl(aq) forms when an > instant before > there was no AgCl(aq) present? Are you saying that the activity > coefficient > of Ag+ decreases sharply when AgCl(aq) forms? If so do you have > any actual > Ag+ measurements to support this view? > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx I think I mislead you by mentioning AgCl(aq) when I meant the dissolved form of AgCl, Ag+ Cl-. What happens is, as you mention, as Cl- concentration ([Cl-]) increases the [Ag+] falls until it reaches a low at about [Cl-]= 0.1M - 1M. As the [Cl-] reaches 1M and above, the [Ag+] rises once more to eventually reach a concentration slightly below the maximum [Ag+] that a very low [Cl-] allows. This is explained by factoring in complexes other than the standard AgCl. > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx > Ivan: As I said, I don't disagree. I just wonder how one could > perform such > an experiment in the laboratory. Perhaps using some buffering agent that > maintains the concentration of Hg(2)2+ constant may work. Can you > propose an > actual experiment? > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx Other than electrolysis of Hg in water, I'm afraid not. > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx > Ivan: Perhaps I'm missing something here, but if you put a little > salt in DW > and apply an EMF similar to that used to generate CS, wouldn't > you get the > following anodic and cathodic reactions? > > Anodic Reaction: [OH-] = 1/2H2O + 1/4O2(g) + 1e Eo = -0.40 > Cathodic Reaction: [H+] + 1e = 1/2H2(g) > Eo = +0.83 > > Net result: No net pH change as water is lost as hydrogen at the > cathode, and > oxygen at the anode. > > Now apply this same electrochemistry to the electrolysis of silver in DW. > > (1)Anodic Reaction (1) Ag(s) = [Ag+] + 1e > > Eo = -0.80 > (2)Anodic Reaction (2) [OH-] = 1/2H2O + 1/4O2(g) + 1e > Eo = > -0.40 > > (3)Cathodic Reaction (1) [H+] + 1e = 1/2H2(g) > > Eo = +0.83 > (4)Cathodic Reaction (2) [Ag+] +1e = Ag(s) > > Eo = +0.80 > > > As I recall, the half cell reaction that occurs first, is the one > with the > lowest E (which equals Eo + the overvoltage to drive the electrochemical > reaction to the right) Agreed >To my way of thinking, the only reason Reaction(2) > above MAY not go to the right is because the overvoltage is so > high that it > significantly exceeds the -0.80 (+ its overvoltage), the voltage > associated > with the oxidation of metallic silver to form silver ions. > I don't see why it > is important to consider what is going on at the other electrode. > This type > of analysis looks at the electrochemical potential at EACH HALF > CELL in terms > of "E" to decide which reaction(s) is most likely considering that the > electrolysis of water will eventually -- if not immediately -- limit each > half cell voltage . So please tell me in terms of "E" why any of > the above > reactions will not go to the right. > xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx The over-voltage of oxygen at silver electrode is about -0.6V + (-0.4)= -1.2V and is unlikely to procede with any vigour while silver is oxidised at -0.8V. It is only valid to look at each half reaction in isolation, if the each is in its own cell, connected by a salt bridge. In this case, while it is convenient to work out the details of half reactions, they must be summed at the end to find the state of the whole cell. As the reduction potential of silver and hydrogen are nearly the same, I suppose that there is a possibility that oxygen could be evolved at the anode whilst silver ions are reduced at the cathode, thus resulting in a rising H+ level. > Roger > Ivan. -- The silver-list is a moderated forum for discussion of colloidal silver. To join or quit silver-list or silver-digest send an e-mail message to: [email protected] -or- [email protected] with the word subscribe or unsubscribe in the SUBJECT line. To post, address your message to: [email protected] Silver-list archive: http://escribe.com/health/thesilverlist/index.html List maintainer: Mike Devour <[email protected]>
