Comon

big314mp, we do. We have three scenarios/conditions in our investigation that prevents the oxidation of shiny metallic iron. 1. iron in NaOH solution 2. iron in freshly boiled distilled water and "sealed" with a petroleum jelly cap 3. iron in non-acidified dichromate solution We setted up three other scenarios to predict and then confirm corrosion of the nails. 4. iron in 0.1M HCl solution 5. iron in a 0.1M neutral salt solution 6. iron in distilled water(with naturally dissolved air in it) The results/observations are "great" - including the much later addition of ferricyanide solution to confirm or refute the presence of Fe2+. The discussion that has gained momentum is centred around the question: Why doesn't an oxidising agent like dichromate (or permanganate or even others) reduce in the aqueous presence of a reducing agent( such as iron) if hydrogen ions aren't present in the solution?

Dear Sotiacho, I have had an opportunity to discuss the matter with a colleague and think about it as well. While the following answer stares at us in the face, I accept that it is simplistic, perhaps wrong and it also raises more questions. If dichromate is going to reduce then it would be in proximity to a reducing agent. However, for the reduction to be successful the electron transfer would have to occur (@ 3e-s per chromium atom) and the ion would have to break up fto result free chromium III ions. (presumably due to a newfound instability). This break-up would be facilitated and occur only if there are a heap of free protons(hydrogen ions) in solution whereby the "freed" oxide ions could bond with two protons to form water. Otherwise there would be free oxide ions in solution and that just doesn't happen in aqueous solution to any appreciable extent; Clumsy, simplistic and perhaps incorrect - but I find it logical for the moment. The main question that is raised is: Does dichromate become unstable when it manages to accept 6 electrons? I guess so, but does the instability only cause the break up to be complete when there are hydrogen ions present to facilitate the boding to oxides as they're released? See you tomorrow.

I'm treating it as a single compound; I'm just showing oxidation numbers for most of the critical species in the overall equation.

Thanks Miguel. PD means potential difference to me i.e. voltage. I will access my radiotron wednesday; i'll look then.

Hi I have a CRT radiotron tube with a fluorescent screen that's been used for electron beam deflection demonstrations in the past. There are four pins and one anode at the top of the tube.(Rated 6.3 V for cathode heating and 500 V max) Haven't been able to get it to operate. Suspect it may be blown, but I wouldn't mind locating any literature on the deflecting electron beam demo with magnet, mainly to get the wiring of the pins right. I do get the glow from the heating of the element at 6.3 V. My power supply also provides 300 V. Is the 300V enough to supply a PD for the beam to flow when the "tron"is rated at 500V max across that gap. Any comments? Any literature? Thanks in advance

Would CuFeS2 dust, that comes off the trucks, react with slightly most and acidic soils to evolve H2S? I'm interested in how CuFeS2 can be oxidised (S in a moist and slightly acidic environment (from 2- to 6+ ?) I'm not sure of the half equations. however for S2 to SO4, S24-(s) +8H2O(l) → 2SO42-(aq) +16H+(aq) +16e- and 2H2O +2e- → H2(g) + 2OH-(aq) So overall, Cu2+Fe2+S24-(s) +24H2O(l) → Cu2+Fe2+2SO42-(aq) +16H+(aq) +8H2(g) + 16OH-(aq) With H+ not being mopped up completely H2S may be evolved in acidic moist soils, right? Hypothesis: Cu2+ free ions enter the soil in moist conditions and H2S will be evolved if the soils are moist and acidic.

Insane Alien, I don't have any access to the ore but I know that metal sulfides evolve H2S in acidic conditions. Is that too simplistic a view?

What could change CuFeS2 to CuFeSO4 in the environment, and how could I determine both of their solubilities please?

Hello all. Some friends of mine are concerned about a future potential mining operation in their forest and near their country town via copper iron sulfide. This ore is going to be extracted and trucked through the town and pristine forest. The concern stems from someone's assertion that CuFeS2 can be oxidised to CuFeSO4. I'm concerned that in acidic moist air, hydrogen sulfide may evolve. I don't understand this in terms of oxidation numbers because it seems to me that copper's oxidation number doesn't change in going from the CuFeS2 to CuFeSO4. Is it because of the aquisition of oxygen that copper iron sulfide is considered to oxidise into copper iron sulfate? More importantly what enviromental chemical issues would be created by the introduction of the CuFeS2 as airborne dust and spillage and the dissolvable "oxidised form" of CuFeSO4 into a forest and nearby township? Any useful comments, chemically based or in terms of consequences to a forest (eucalypt)environment and a township of people through which the trucks will roll would be appreciated. Thank you again.