# Precipitating Metallics.



## CattMurry (Dec 15, 2021)

Hey there my expert friends! My question today is discussing the formation of matallics! I had a chat with the youtuber OwlTech who was very helpful about the subject, and he had said that precipitation had to do with negative charged compounds which take the added electron from the aqueous "metal" solution which has a positive charge. My question would be if this is the case SMB has a valency charge of (-2) with the bisulfite being the main reactant, so would sodium bicarbonate work as well being also a (-2) valency charged compound? I actually saw OwlTech use it after I went on the research escapade, but he has also used sodium hydroxide, and fomic acid, but I am unsure why. He has said this: "it's the CO that does the trick from decomposition of fomic acid", but carbon monoxide (CO) has a zero ionic charge. My main question would be can sodium carbonate replace SMB for precipitation, or even be able to deNOx a solution replacing urea, and other compounds? Also, what else is at play for precipitation other than ionic charge relationships? Thanks everyone for your answers to my other questions so far, and for this thread as well! A very helpful, polite, and informative community.


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## Lino1406 (Dec 15, 2021)

S of bisulphite or SO2 is +4, going to +6 on reduction (or precipitation) hence action is -2. Required also correct potential between reductant and oxidizer.
CO2 does not have that ability.
Formic acid can also turn to CO2, thus giving 2 electrons


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## CattMurry (Dec 15, 2021)

Lino1406 said:


> S of bisulphite or SO2 is +4, going to +6 on reduction (or precipitation) hence action is -2. Required also correct potential between reductant and oxidizer.
> CO2 does not have that ability.
> Formic acid can also turn to CO2, thus giving 2 electrons


If I understand correctly you have said that the SO2, and SMB have the ionic action of (-2) due to them giving the 2 electrons? but the formic acid/CO2 does not have that action? Just a little confused on the last part there of: "thus giving 2 electrons". Wouldn't the formic/CO2 have the same tradeoff, or ionic action? Thanks for the answer. Just wanted to also ask about the Sodium Carbonate's action as well, and thanks for the patience if I have this backwards haha.


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## butcher (Dec 15, 2021)

We can give electrons back to an ionic compound to change its valence or even reduce the ionic salt back to being an elemental metal, or even change the type of ionic salt the metal ions is (say from being a metal chloride salt into becoming a metal sulfate salt where one metal may precipitate from solution and a different metal may not become insoluble), by the use of many different chemicals or metals.

Not all metal ions react the same to a change in valence or a to a change in the type of salt it is, or to the ionic salt that it can form, some metals may precipitate and others may not in an adjustment of their valence, or its type of ionic compound from the same solution.

We normally make adjustments with _certain_ chemicals or metals or use certain reducing agents to selectively remove or change certain metal ions from solution while leaving others in solution, like removing lead from solution while leaving the gold in soluble form, even if both metal ions valence has been changed or they both can change the types of salts they are or could become.

But our goal may not be to just change the valence of the solution, with just any chemical or metal that we can use. We may have several different metal ions in solution at once, and we do not wish to change them all at once or reduce them or precipitate large groups of different metals at once.

Our goal may not be to reduce all metals in solution back to elemental form, or to precipitate most anything, or just anything from the solution, or to convert large groups of different metal ions to a carbonate, hydroxide, or oxide or to some other elemental form.

We are picky refiners, and also tend to be selective in what metal or metals we wish to convert or change in form or valence.

We want to separate the different metals from solution selectively and cleanly as possible as much as we can we try to which metals we reduce from solution and which ones we don't, which chemical or metal that will be more selective for the ions we are after, or where one metal may precipitate by changing the valence...


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## Lino1406 (Dec 15, 2021)

Yes formic acid/CO2 route has same action, not the CO route. Carbonate is acting as Na2O + CO2 nor Na2O neither CO2 has same action


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## CattMurry (Dec 15, 2021)

Lino1406 said:


> Yes formic acid/CO2 route has same action, not the CO route. Carbonate is acting as Na2O + CO2 nor Na2O neither CO2 has same action


Okay, so him stating that the monoxide (CO) doing something is irrelevant then? While CO2/Formic does have the same reaction, but it is not the appropriate action required? What do you mean by the sodium oxide, and carbonate not having the same action? Is that in relation to each other, or comparing their reaction to the stated reactions above? The youtuber just explained it as a ionic action of (-2) which cements matallics due to them getting their electrons back. In such case I would then assume any compound like baking soda (sodium carbonate) which does have that ionic action of (-2) to be sufficient from what was said. I was curious if there was anything else to it than just that, or if it truly just requires that (-2) action. Thanks for the replies!


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## CattMurry (Dec 15, 2021)

butcher said:


> We can give electrons back to an ionic compound to change its valence or even reduce the ionic salt back to being an elemental metal, or even change the type of ionic salt the metal ions is (say from being a metal chloride salt into becoming a metal sulfate salt where one metal may precipitate from solution and a different metal may not become insoluble), by the use of many different chemicals or metals.
> 
> Not all metal ions react the same to a change in valence or a to a change in the type of salt it is, or to the ionic salt that it can form, some metals may precipitate and others may not in an adjustment of their valence, or its type of ionic compound from the same solution.
> 
> ...


Not quite sure what this explains, but thanks for the reply nonetheless! I am aware of pickiness, and the process of creating salts, or other changes to maneuver around the aqueous compounds to then drop, or react what you want. Such example that is clear is how some people choose to put everything (precious metals) into solution, then use the ionic action of HCL (-1) ionic charge to bump, or precipitate silver chloride (0) ionic charge from it's silver nitrate aqueous (-1) ionic charge if I'm not mistaken. After that you give another electron to the chloride salts to make Silver Oxide with a ionic charge back to (+1) with hydroxide, or other compounds.


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## FrugalRefiner (Dec 15, 2021)

You're confusing two different principles.

When we precipitate gold from a solution, we usually do so by "reducing" it. Gold has been oxidized to put it into solution. As it is oxidized it gives up electrons and it changes from a solid metal to an ion in solution. 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.

When we precipitate silver from a solution, we can either reduce it or we can create a nearly insoluble compound that precipitates. If we use copper to cement the silver from solution we are reducing it. The solid, metallic copper gives up electrons to the ionic silver that is in solution. The silver is reduced and precipitates as a solid metal as the copper is oxidized and becomes an ion in the solution, replacing the silver. But if we add HCl or table salt (NaCl), the chloride ions combine with the silver ions to form silver chloride (AgCl), which is nearly insoluble and the AgCl precipitates.

In all these cases, we have a precipitate that collects on the bottom of our beakers. But in some cases the precipitate is a solid metal while in others it is a compound of our target metal combined with another element or ion.

Dave


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## CattMurry (Dec 15, 2021)

FrugalRefiner said:


> You're confusing two different principles.
> 
> When we precipitate gold from a solution, we usually do so by "reducing" it. Gold has been oxidized to put it into solution. As it is oxidized it gives up electrons and it changes from a solid metal to an ion in solution. 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.
> 
> ...


Good explanation Dave. Side reactions, or many different intermediates can make it hard to keep track of haha. Looks like silver chloride does have some decent solubility in ammonia, and even apparently hitting it with light will bring it back to the oxide. Fascinating stuff. Still a little confused on why baking soda isn't just thrown at everything if it does give electrons as well. For the rest of understanding how some metals dissolve, and precipitate over other metals first the reactivity metals chart says it all, but of course it is not so clear cut in many situations, and practice. (if any of this is inaccurate do correct! I am no chemist haha.)


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## butcher (Dec 15, 2021)

Still a little confused on why baking soda isn't just thrown at everything if it does give electrons as well.

We are not trying to convert everything carbonates or oxides.
The reactivity series does not explain it all, although it is a very helpful tool like our others to give some explanation.

some metal ions will precipitate with an alkaline substance (like ammonia or hydroxide...) and then redissolve in excess of that same alkaline solution... Look into the amphoteric metals, although gold is one of those metals, although the gold ions would have to be subjected to more extreme conditions than many of the more reactive metals.

it is about electrons, either lose or gain of, or of sharing or trading places in ionic solutions,

There are many factors that can come into play, Not just valence. and many different conditions that can affect results on the chemistry of the metals or their ionic compounds. valence is not all there is to it.


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## orvi (Dec 15, 2021)

CattMurry said:


> ... he had said that precipitation had to do with negative charged compounds which take the added electron from the aqueous "metal" solution which has a positive charge. My question would be if this is the case SMB has a valency charge of (-2) with the bisulfite being the main reactant, so would sodium bicarbonate work as well being also a (-2) valency charged compound? I actually saw OwlTech use it after I went on the research escapade, but he has also used sodium hydroxide, and fomic acid, but I am unsure why. He has said this: "it's the CO that does the trick from decomposition of fomic acid", but carbon monoxide (CO) has a zero ionic charge.


I don´t think we can judge this so easily, but i could give my bit of experience with theory here.

The assumption with "negatively charged" particles... I thing i get somehow what it should mean: _metal ion in solution is positively charged (cation). in order to get it back to the metal, it must come to the contact with some "negatively charged" particle, which give the electrons to it, make it precipitate as non-charged metal._
I will not consider it this way, but in very simple (and clearly somewhat misleading) way, OK. Could be used to disclose chemistry in "friendly way" for non-chemists  

Thing is, as it was stated by more of us here, compound which cause metal cation to go back to the non-charged metallic form need not to be charged. It just need to have some electrons to offer 
Here we come to thing called "redox potential". I think the best approach to disclose what is happening in your pot.

*All particles/chemicals in your solution want to get in the form, where they would have the lowest energy possible.* There are chemicals that very kindly pass electrons, and there are chemicals that very kindly accept some. Ones are in our eyes "reducing agents" and the others "oxidating agents". Oxidants accept electrons, reducers pass electrons.

This is when the "redox potential" comes into the game. Redox describes certain half-reactions in volts, compared to so called "standard hydrogen electrode". Not exactly needed to explain what it precisely is, but it is our "zero potential".
As we mix together two chemicals, say gold chloride and sodium sulfite, they will interact together.
We could describe this in two half-reactions:
-reduction of Au3+ to metallic gold: Au3+ + 3 electrons ---> Au(0)
-oxidation of sulfite to sulfate anion: SO3- + H2O - 2 electrons --->HSO4- + H+
We focus only on the electron transfer reactions. Everything else is ommited for clarity. Sodium and chloride ions combine to make NaCl, but really they dont gain or loose electrons, so they dont figure in half reactions.

Both have they redox potential. Half-reaction with gold has value of +1,40 V. Half reaction for sulfur dioxide has value of +0,40 V (these values differ as temperature, concentration and other ions are taken in consideration). 
When combined to give result: oxidant minus reductant. +1,40 - 0,40 = +1,0 V
Value is positive. This indicates, that reaction will proceed.

If we take copper for example:
Cu2+ + 2 electrons ---> Cu(0), this half-reaction has potential of +0,34 V.
If we try to combine this copper half reaction with above sulfite reaction, we will find that +0,34 - 0,4 = -0,06 V
Value is negative, so reaction will not proceed.

This is theory. In real world, ions are complexed with other ligands (AuCl3 in solution likely form AuCl4- = lowering the energy), temperature, concentration and other factors hassle with these ideal conditions numbers.

But this is what is happening. It could be considered as some kind of trade. 
When somebody try to sell you some stuff, you could accept or reject the offer. The same happens with chemicals in your pot. Copper in solution rejects the sulfite offer, but gold kindly accept it


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## GuyGuyGuythe1st (Dec 15, 2021)

Thank you all for the knowledge silo. Very informative and helpful.


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## CattMurry (Dec 15, 2021)

orvi said:


> I don´t think we can judge this so easily, but i could give my bit of experience with theory here.
> 
> The assumption with "negatively charged" particles... I thing i get somehow what it should mean: _metal ion in solution is positively charged (cation). in order to get it back to the metal, it must come to the contact with some "negatively charged" particle, which give the electrons to it, make it precipitate as non-charged metal._
> I will not consider it this way, but in very simple (and clearly somewhat misleading) way, OK. Could be used to disclose chemistry in "friendly way" for non-chemists
> ...


Nice in depth explanation my friend. This was the part I had a hard time finding where to even start looking:

-"Thing is, as it was stated by more of us here, compound which cause metal cation to go back to the non-charged metallic form need not to be charged. It just need to have some electrons to offer.
Here we come to thing called "redox potential". I think the best approach to disclose what is happening in your pot."

Then:
-"This is when the "redox potential" comes into the game. Redox describes certain half-reactions in volts, compared to so called "standard hydrogen electrode". Not exactly needed to explain what it precisely is, but it is our "zero potential".

This is why I came here. I couldn't understand how it could be that easy as in (just finding - charged compounds). It all has a lot of arms to it due to the redox potential, valency, How much of what is in "the pot", temperatures, environment, or atmosphere, and I'm sure a bit more. Apologies to the other users for not having those facts stick to my understanding due to my ignorance! I need basically all the details at once to even start understanding this complex web. This will help clarify, but do you know where I can learn a bit more on that half reaction math? Those are some basic, but wonderfully put examples for my potato head lmao. Thanks again for this!


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## CattMurry (Dec 15, 2021)

butcher said:


> Still a little confused on why baking soda isn't just thrown at everything if it does give electrons as well.
> 
> We are not trying to convert everything carbonates or oxides.
> The reactivity series does not explain it all, although it is a very helpful tool like our others to give some explanation.
> ...


I understand... I think. For your example there is a balance, or goldie locks zone for many things, and when the acid, or alkaline reactants get too concentrated it then becomes resoluble into the solution, and even vise versa. I did not mean to say the metals reactivity chart is the golden tool of knowledge for this either haha.


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## butcher (Dec 15, 2021)

We use many different tools to try and describe what could happen or try and describe what is happening, or to get some kind of guess why something is not happening that we expected to, valence change, is only one of the tools.

Things like temperatures boiling points, acidity or pH, concentration as well as other factors also come into these figures of how reactions may proceed,

Exchange of metal ions, electron movement, potential difference, and general chemistry, valence is only one tool we have many more tools in the toolbox,.

Solubility, KSP and or the solubility rules, ORP, oxidation or reduction potentials, redox reactions, chemical and physical reactions, oxidants and reducing agents, selectivity, catalytic actions, anion and cation exchange, rationalizations, precipitations, crystallizations writing and balancing chemical reactions or equations, vaporizations, charge or dissociation of electrically charged ions, ionic exchange, specific gravity, electron shells of different metals and the movement of electrons, molar concentration, percent hydrogen, atomic weights, and chemical conversion...


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## CattMurry (Dec 15, 2021)

butcher said:


> We use many different tools to try and describe what could happen or try and describe what is happening, or to get some kind of guess why something is not happening that we expected to, valence change, is only one of the tools.
> 
> Things like temperatures boiling points, acidity or pH, concentration as well as other factors also come into these figures of how reactions may proceed,
> 
> ...


I'm aware of this, but thank you for the full scope of vocabulary, and reiteration.


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## butcher (Dec 15, 2021)

Well, i am having trouble explaining and you apparently do not understand what I am trying to say.


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## CattMurry (Dec 16, 2021)

butcher said:


> Well, i am having trouble explaining and you apparently do not understand what I am trying to say.


What are you having issue explaining? I simply only wondered if the youtuber was being honest which he was, but his explanation is quite a bit simplified. None of you really answer strait forward the main question either. Baking soda is a reducer, and it neutralizes many acids that I can think of just like SMB, and others. Obviously you wouldn't just throw it into a melting pot of solutions lol. You would for example make peracetic acid solution to put all reactive metals like lead, tin, zinc, and so on into solution. Maybe a H2O, then HCL washing to make sure there's no random crud around (of course Cu will resist this step if present). Then knowing what else you have in here (lets assume silver, gold,) it would be nitric to eat the silver into its nitrate. Decant, or filter. Precipitate Silver aqueous with HCL, or salt, then in your other container (Au pot) you could go AP route, Chlorine route, or aqua Regia route. Then precipitate the Au with SMB after PH is controlled (in all these instances of precipitation this would be the case for controlling PH). I was simply asking why can't Sodium carbonate reduce it out if it is a reducer. I got every obvious answer but my question really. Lol. So what is hard to explain, and what exactly am I not understanding in all of this? Of course I do not know libraries of knowledge on every reaction, but I realize you wouldn't reduce everything at once, and PH of solution matters along with temperature, and all those other things.


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## CattMurry (Dec 16, 2021)

butcher said:


> Well, i am having trouble explaining and you apparently do not understand what I am trying to say.


Basically what I am looking for is the knowledge that points to the action, or paper describing why hydroxides are preferred, or things like SMB, salt, even HCL, and many other's for dropping metals from solutions. I have seen combinations with baking soda, but not just itself used. Was just curious on Sodium carbonates reactions, uses, it's downfalls, and it's potential, or non-potential, and why basically. In my head yes to dissolve things there are many routs to take, but to get the actual oxides, or metallics.... they just end up as metallics. Meaning it seems very strait forward if you have of course the right conditions in your solution. Just wondering on that, because the dropping action seems like there are many different ways to just form the metals back to their metallic state, but it always ends with their matallic state. Rather than the metals going into solution which takes many forms. So, again. I am curious on reducers, and what action they are all doing to do the same thing, and why some are used at the price they are.

*I should add that some reducing agents are bases, and some are acids for obvious reasons. Such example is chlorine rout. Obviously you aren't looking to eat your beaker and adding more caustic/alkaline substances which will for the most part not help in any way. Maybe they can due to what was explained by making a solution too basic/acidic, but that's what I am asking about as well. Where is this math, papers, or diagram for this "goldie locks" zone as I had put it.


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## CattMurry (Dec 16, 2021)

orvi said:


> I don´t think we can judge this so easily, but i could give my bit of experience with theory here.
> 
> The assumption with "negatively charged" particles... I thing i get somehow what it should mean: _metal ion in solution is positively charged (cation). in order to get it back to the metal, it must come to the contact with some "negatively charged" particle, which give the electrons to it, make it precipitate as non-charged metal._
> I will not consider it this way, but in very simple (and clearly somewhat misleading) way, OK. Could be used to disclose chemistry in "friendly way" for non-chemists
> ...


I also found what you were saying when it was in my face the entire time. SMB is an example of what you were referring to:
"Thing is, as it was stated by more of us here, compound which cause metal cation to go back to the non-charged metallic form need not to be charged."
I did not notice that SMB is at 0 charge itself, and that its reaction into SO2 is actually what is doing the "dropping/cementing/etc". I've seen SO2 bubbled into solutions, and it looked heavily ineffective, or just takes forever, so I didn't even think it would be the gas. "SO2 cements out the metals with sodium ions", is what one user on here (website) from another post had said, but how is this if it is only made of Sulfur, and oxygen? I just get to this: *Sulfur dioxide* can bind to metal ions as a ligand to form metal sulfur dioxide complexes, typically where the transition metal is in oxidation state 0 or +1. So, is it the sulfur itself, or as one user had said in another post Na binding ? Also, I still can't find for the life of me redox action potential of different compounds. Ah, standard reduction table. Found it. Of course at 25c, but how can I see different tables for heat? 

Now my head hurts. lmao.


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## Geo (Dec 16, 2021)

Ion exchange - Wikipedia







en.wikipedia.org


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## CattMurry (Dec 16, 2021)

Lino1406 said:


> S of bisulphite or SO2 is +4, going to +6 on reduction (or precipitation) hence action is -2. Required also correct potential between reductant and oxidizer.
> CO2 does not have that ability.
> Formic acid can also turn to CO2, thus giving 2 electrons


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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## CattMurry (Dec 16, 2021)

Geo said:


> Ion exchange - Wikipedia
> 
> 
> 
> ...


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.


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## CattMurry (Dec 16, 2021)

CattMurry said:


> 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.


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 (Dec 17, 2021)

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.


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## CattMurry (Dec 17, 2021)

Geo said:


> 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.


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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## CattMurry (Dec 17, 2021)

Geo said:


> Ion exchange - Wikipedia
> 
> 
> 
> ...


I also read this a bit right now, and it does help a lot with understanding charges. Thanks friend.


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## olawlor (Dec 17, 2021)

Geo said:


> CO2 is a reducing agent for copper. In smelting, carbon is added to reduce copper...


I'd say carbon or carbon monoxide is the reducing agent.

CO2 is the oxidized product of that reduction.

(You can't reduce metals by pumping hot CO2 through them, but you can reduce them by pumping in hot CO.)


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## Geo (Dec 18, 2021)

olawlor said:


> I'd say carbon or carbon monoxide is the reducing agent.
> 
> CO2 is the oxidized product of that reduction.
> 
> (You can't reduce metals by pumping hot CO2 through them, but you can reduce them by pumping in hot CO.)


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.


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## CattMurry (Dec 18, 2021)

olawlor said:


> I'd say carbon or carbon monoxide is the reducing agent.
> 
> CO2 is the oxidized product of that reduction.
> 
> (You can't reduce metals by pumping hot CO2 through them, but you can reduce them by pumping in hot CO.)


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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## CattMurry (Dec 18, 2021)

Geo said:


> 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.


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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## CattMurry (Dec 18, 2021)

Geo said:


> 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.


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.


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## orvi (Dec 18, 2021)

CattMurry said:


> 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.


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## CattMurry (Dec 18, 2021)

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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