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Silver compounds such as silver fluoride are light sensitive (as are all other silver halides) so they will decompose when exposed to light.

I'm stuck with a problem : I want to expose silver(I)fluoride to light to then produce elemental silver and fluorine gas - I'm worried that because of the high reactivity of fluorine silver(I)fluoride will be formed again...

So how could I successfully split silver(I)fluoride by light into elemental silver and fluorine gas without the fluorine reacting again with the silver?

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    $\begingroup$ If you want to actually keep the fluorine, you could pull the fluorine gas away with a vacuum and have it pass through a liquid nitrogen trap or two (fluorine boils only 8 K above LN2, however, so you'd need a system with excellent heat exchange). Though liquid fluorine is probably not something I would ever like to see without some serious PPE. It could also damage the pump irreversibly. If for some reason you want to split AgF and not keep the fluorine, it's pretty easy, pretty much any substance will react with fluorine gas preferentially to silver, so you can put some Al foil nearby. $\endgroup$ – Nicolau Saker Neto Dec 11 '13 at 0:00
  • $\begingroup$ and could teflon be used (or any other fluorine-containing 'polymer') to transport the fluorine gas away? (I suspect that it won't react with fluorine?) $\endgroup$ – user2117 Dec 11 '13 at 12:36
  • $\begingroup$ Teflon or other saturated perfluorinated polymers such as FEP are indeed resistant to corrosion by fluorine. I don't know how they behave at cryogenic temperatures, though. Do you have the proper equipment and training to handle fluorine? If not, don't try to isolate it. Many people have died or been severely poisoned in the process. $\endgroup$ – Nicolau Saker Neto Dec 11 '13 at 19:13
  • $\begingroup$ Actually scratch that, what I suggested won't work. Under low pressure, fluorine's boiling point will definitely be lowered below nitrogen's, so it won't condense at all. I should've realized that sooner. It's probably best if you find some references for this sort of stuff, though I don't know where you might find them. $\endgroup$ – Nicolau Saker Neto Dec 11 '13 at 19:24
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    $\begingroup$ No idea. You should be able to get a feel for the spontaneity of the reaction by drawing a thermodynamic cycle and using the Gibbs' free energy state function properties, but I don't know how to obtain kinetic information. The reaction speed may simply be limited by the incident photon flux with sufficient energy to drive the reaction, though, and in that approximation all you have to do to increase speed is find a more powerful source of light of the right wavelengths. $\endgroup$ – Nicolau Saker Neto Dec 13 '13 at 13:20
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As discussed in the comments, the asker only wanted to create some fluorine in situ for use in another flask, with no desire to store it. Chemistry with fluorine is difficult due to its dangerously high reactivity. Do not attempt anything without the proper safety procedures.

I actually still have doubts whether it is possible to obtain fluorine gas reliably from the exposure of silver fluoride to sufficiently high energy photons, and if there is any simple way to produce photons with the required energy. It is well known that the other silver halides easily suffer photolysis, especially for the heavier halides, but reactions which produce fluorine gas tend to be extremely endergonic, and therefore don't happen except in rather severe situations.

In case it is somehow possible to produce fluorine gas directly by photolysis of silver fluoride, then it is possible to transport the gas to another container where it can react, simply by connecting the flasks with a pipe. Few materials are capable of withstanding exposure to fluorine, however, so the containers must all be made of very well dried gas. Unless the asker has access to a very high power light source, I suspect that the rate at which silver halides decompose is quite slow, which may preclude their efficient usage as halogen generators. This is especially true for silver fluoride, which likely decomposes the slowest, and the fluorine formed in very low concentrations could easily be lost by reaction with impurities or residual water before even reaching the target. I do not know whether there are any catalysts which can be added to increase the photolysis speed of silver fluoride. Presumably some semicondutor with a high bandgap, such as titanium dioxide, could absorb the light and cause photoreduction of the silver/photooxidation of fluoride, but the catalyst would be quickly consumed by the fluorine formed.

If fluorine gas can be formed by silver halide photolysis fast enough, then by simply linking it to another flask with the desired reactants should cause the fluorine to diffuse and react, followed by a continuous equilibrium shift in the production of more fluorine in the flask containing the silver fluoride, and its consumption in the flask with the other reagents.

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