Precipitating Metallics.
- Thread starter CattMurry
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You had the most straight forward answer, so thanks again. I got and was also confused on chemistries perfectly odd terminologies. Example:
Oxidizer- to lose electrons, and gets (+) action,
Reducer- To gain E, and gets (-) action
Reducing in English is normally a word for loosing something, or getting smaller
Then what Dave said: "A reducing agent, like SMB or ferrous sulfate, gives up electrons to the gold, which then changes from an ion back to a solid metal."
Then:
A reducing agent, like SMB or ferrous sulfate, gives up electrons to the gold, which then changes from an ion back to a solid metal.
Yet SMB is reduced itself, and gets a oxidizing action to then make SO2 release.
-reducers have extra electrons (that is, they themselves are reduced)
So then it looks all backwards because they explain it like Oxidizing, and reducing is the action going to happen while that is not the case (I can't even tell anymore though).
In reality reducers are reduced themselves having more E's already, and get oxidized
While Oxidizers are already in oxidized form, and then get reduced.
BUT at the same time there are layers or levels to those 2 forms lmao.
My thought was: If the ions in solution have lost E in the oxidizing action how does a reducer which gains even more E cement the ionic metals... but no. Ionic solution that's oxidized gets reduced by a reduction action, but that action looks like this.
SMB breaks down into SO2 which looks to be oxidation, Then SO2 with its +4 oxidation state can do it's work which in this case is undergoing it's own oxidation to reduce, or give the ionic metals E's. Seems this is why SO2 is popular Sulfur dioxide is also a good reductant, as well as an oxidizing agent. BUT then again here is CO2- Some of the carbon atoms in CO(g) are oxidized from +2 to +4, and some of them are reduced from +2 to 0. Thus carbon atoms in CO are both oxidized and reduced, and CO is both the oxidizing agent and the reducing agent. So I still don't get what type of CO2, or any gas makes it appropriate to be a reductant.
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I need to see potentials I think.
like here: Standard electrode potential (data page) - Wikipedia
but I am new to this, so putting this into actual practice will take a good while lol.
This part I butchered. I can't even tell if what google told me was Co (colbalt), or CO (carbon monoxide). It appears that CO2 is carbons oxidation state of +4 which is C's max, thus It cannot be reduced. I am confused why it cannot be oxidized, and reduce other ionic metals out of solution though. SO2 is used instead with an oxidation number of also +4 on the S, and CO is at +2 on the C. I still can't figure out why CO2 wouldn't be used, or doesn't work.... Would it not be oxidized back to CO? Or get back its electrons in a reduction action loosing an oxygen, turning into CO, then in essence oxidize the metals?
Here is another chemistry trolling vocabulary example. CO2 can't be reduced? Yet it is explained that it gains or frees up its electrons by going from CO2+4 into CO+2, which looks like a reduction it seems, yet it has more electrons in CO due to not being bonded to another O. Then oxidations are also explained as loosing electrons, but also it is gaining an O for bonding, which is actually loosing electrons that are free to bond on C. Not understanding terminology over here much it seems that chemistry keeps flip flopping what it is explaining, then how it actually plays into it's partnering to other elements.
Then google:
Is CO2 an oxidizing agent?
The abundant availability, non-toxic, economic and mild oxidizing properties of CO2 has resulted in immense interest in its use as an oxidant in several reactions, such as the oxidative coupling of CH4 and the oxidative dehydrogenation of alkanes and alkyl aromatics.
To oxidize another thing it must be reduced. Why does google say CO2 can't be reduced, and that it can oxidize?
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Geo
Well-known member
CO2 is a reducing agent for copper. In smelting, carbon is added to reduce copper and any time you have metallic copper, gold will not stay in solution, even if the solution is a couple thousand °F. It also reduces iron in a bloomery forge. The earliest bloomeries was a trench dug into the side of a hill. Dried trees were laid in the bottom and iron containing ore was placed on top of the wood with more wood added in layers. Then a last layer of wood was added and the whole thing covered with earth. The layers were not compacted as it needed air flow from bottom to top. Then they lit the bottom and and the flame and heat would travel up and out like a chimney flue. Once it burned completely out, they would dig it open and recover any lumps of iron that had collected. The iron was reduced by CO2 that was formed by the incomplete combustion of the wood. Iron is always mined as an oxide. The extra oxygen was stripped in the production of CO2 and expelled reducing the iron oxide to iron metal.
That's pretty fascinating honestly... Could you elaborate on this first part? Are you saying solution as in the smeltery? Also, how does smelting work with reducing the copper? Just having a hard time envisioning it to be honest. (nvm I forget you are working with oxide in ores mostly in smelting. Thus, carbon's oxidation as a reductant makes sense, but let me know if I have that understanding strait). I was more confused on why SO2 is always used, and why CO2 cannot do what SO2 does in solutions (aq) of metal ions/salts/cation solutions, but I think I might have found it. The only difference looks to be that S had multiple stages of + charges, and a - charge (ionization possibilities) in which C doesn't. looks odd to me, and I don't fully understand it. S has -2, +2, +4, +6 making it very liquid in it's uses it seems. Uck. I really hate sulfur though lmao, but I digress. Thanks for this info. A neat way to oxidize with a very long superheated burn tube in essence.
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Geo
Well-known member
You are right on both points. But, CO2 is a byproduct of removing the oxygen from an oxide. CO is also produced in a bloomery due to incomplete combustion of the wood. CattMurry Was correct in assuming that the production of CO2 was what actually reduces the copper oxide to copper metal. By combining the oxygen to carbon at high temps, the oxygen is removed from the ore. It's not a complete reduction because CO2 is not a reducing agent. The effect is by happenstance and not design. But, if it works, it works. As an aside, super heated CO2 is pumped into steel foundry melting pots to stir the molten metal.
Thanks for the answer. I had actually seen that, and to my very basic noobish eyes it would look like CO2 could be reduced to oxidize (X). (used as an Oxi agent).
When I search I get something like this:
This means that, in a chemical reaction, CO2 can act only as an oxidizing agent, as it occurs, for instance, in the photosynthesis process. Whereas, CO can both increase and decrease the oxidation state of its carbon atom, according to the circumstances and thus acting as both oxidizing and reducing agent
So, yeah. It take some apparent special exchanges that I am right now unaware of to reduce the CO2 to CO, and for it to be used for an oxidizing agent it must meet certain criteria. Makes sense though that CO can be oxidized to CO2 to then reduce (x) though.
So, my very new understanding of everything chemistry is getting mixed up on the redox reactions of metals going to their metallic versions, rather oxidized versions, or even salts, and other's I think. Lol. Got me in quite a twist.
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Just wanted to shoot a quick question for SMB. So, does it make SO, into SO2 in redox, to then reduce (x)? Or does it make SO2 into the colorless SO3 through redox to be able to reduce (x)?
Also, if this just makes no sense, or is literally backwards just say that if you wish to reply haha. Still confused on the versions of aqueous metal solutions, and the forms that the metals are in, then go into a bit.
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Ok I can read this which is interesting to look at:
Tetrachloroaurate Ion + Sulfur Dioxide + Water = Gold + Sulfate Ion + Chloride Ion + Deuteron · FeCl3 + H2CO3 = HCl + Fe2(CO3)3 · Ba(OH)2 + NH4NO3
So, looks like redox is happening between the waters (O) and SO2 to get SO4 while some action is breaking the CL bonds of the salt. That action is what I do not know what to search for basically.
Redox is happening when atoms in the molecules change their oxidation state. Water have oxygen in -II oxidation state. Sulfate also have oxygens with -II oxidation state. No change = no redox.
On the other hand, sulfur in SO2 is in +IV oxidation state, but at the end of reaction, it converts to sulfate, where sulfur is in +VI oxidation state. The oxidation part of the redOX is happening here.
On the other side, chloride ligands on aurate ends up being salt. Chlorine is in -I oxidation state, so no change = no redox. But gold in aurate anion is in +III oxidation state, and at the end of the reaction gold is metallic, thus oxidation state 0. This is the reduction part of the REDox.
Thank you for the technical touch there. Now things are adding up correctly to my eyes lmao. I still didn't get what rules/laws explaining the breaking of the bonds of the Chloride salt's bonds. If I missed the reasoning my apologies. How do the ligands break? (Disassociates I should say)
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