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I was so intrigued by the reported observation that sunlight precipitates the iron from $\ce{K4Fe(CN)6}$ (originally attributed to Matuschek, 1901) that I wanted to see it myself. I put a saturated aqueous solution, uncovered, in direct sunlight for four hours at mid-day. Absolutely nothing precipitated from the solution!

My understanding from the given reference is that in solution the $\ce{K4Fe(CN)6}$ gives $\ce{Fe(CN)6^{4−}}$ and then:

$\ce{Fe(CN)6^{4−} + 2H2O <-> Fe(CN)5 + (H2O)^{-3} + HCN + OH-}$

The photolysis causes $\ce{Fe(CN)6^{4−} ->[h\nu] Fe(CN)6^{3−} + e-}$, so then we also have

$\ce{Fe(CN)6^{3−} + 2H2O <-> Fe(CN)5 + (H2O)^{-2} + HCN + OH-}$

Is it correct to assume that if a $\ce{Fe(CN)6^{4-}}$ anion absorbs a photon of adequate energy (in this case it appears to require $\lambda < 313nm$) then with 100% probability the anion will photolyze the Fe bond?

And are there models that predict (at least order of magnitude) the absorption probability of a photon by an aqueous anion (being irradiated with a given spectral flux)?

Furthermore, what is the reaction that would precipitate iron in this case? When we bring the K back into the equations it looks like the HCN would react to produce KCN. I don't know what should happen to $\ce{Fe(CN)5}$.

In fact, after several days in open sunlight the solution is subjectively getting darker and I do see a small amount of red sediment at the bottom of the dish that appears to redissolve on swirling. Could be an iron oxide, but would also be consistent with $\ce{K3Fe(CN)6}$. In either case I don't have any indication from the source what should happen to the $\ce{(CN)5}$ groups if iron is precipitating, or what else might take one of the K atoms from the initial compound.

(Ultimately I had hoped to quantify the rate of photolysis by weighing the precipitate.)

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  • $\begingroup$ You do realize that the glassware (unless it's quartz) strongly absorbs UV light below 300 nm? If you want to use UV light below 350 nm, I suggest you shine it directly on the solution and not have it pass through glass. $\endgroup$ – LDC3 Jul 26 '15 at 16:41
  • $\begingroup$ @LDC3: Yes, it's a shallow glass container that I left uncovered. Presumably the heating of the solution by the sunlight and elevated ambient temperature is offsetting any evaporation that might be occurring that would cause the initial solute to precipitate. $\endgroup$ – feetwet Jul 26 '15 at 16:49
  • $\begingroup$ If your solution is basic, then $\ce {Fe(OH)_3}$ will precipitate. $\endgroup$ – LDC3 Jul 26 '15 at 20:17
  • $\begingroup$ @LDC3: I confirmed its pH is 8.6. The amount of red sediment that has settled on the bottom is miniscule. Also, curiously, what little is there does not follow a magnet. So if some form of Fe is precipitating but staying in suspension there goes my idea of magnetic separation.... $\endgroup$ – feetwet Jul 29 '15 at 23:05
  • $\begingroup$ 8.6 is barely basic. I was thinking 10 or higher. $\endgroup$ – LDC3 Jul 30 '15 at 0:28
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I would put the solution under a blacklight exposure source or a equivelant such as a mercury arc lamp at close range. The referenced writing makes no specifics about the the timeframe of prolonged. This could have been 3 months of exposure to sunlight.

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  • $\begingroup$ Good point. I'll have to check sub-300nm luminance ratings on shortwave UV lamps and compare with peak daylight. I had a short period in mind because on the next page of the reference in my question it notes that a recent paper (which I've requested but not yet received) showed complete photolysis by sunlight within 30 minutes, but that was of a miniscule concentration of $\cf{Fe(CN)6^{-4}}$.... $\endgroup$ – feetwet Jul 26 '15 at 1:00

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