RogueRose
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Single displacement reaction with metal salts
So if you want to displace the copper in copper salts by precipitating the copper metal by placing a more reactive metal in the salt solution will
this work with any metal that is more reactive or does there need to be a specific "difference" in the reactivity of the metal?
Now I know that there isn't the same "distance" in the reactivity, like a continual numbered line. If the difference in reactivity were numbered it
might look like this where the # is the distance to the next element.
A .5
B .02
C .85
D 1.6
E .14
F .36
So A-F would be a total of 3.11 and A-C would be .52.
So here B->C would be very small difference where D->E would be about 80x greater distance than B->C. So in this example C might not
displace B in a salt solution.
What I want to figure out is how can I determine what elements will replace another if I'm using a solution of either a chloride or sulfate (I guess
nitrates would work as well..?)
This list is more reactive as you go down the list
Silver
Mercury
Tungsten
Copper
Bismuth
Antimony
Lead
Tin
Nickel
Cobalt
Cadmium
Iron
Chromium
Zinc
Manganese
Titanium
Aluminum
Beryrillium
Magnesium
calcium
Barium
Lithium
Sodium
Potassium
I'm wondering if I put antimony in either CuSO4 or CuCl2, would it dissolve the Sb and precipitate Cu?
Also would Zn replace Fe in FeSO4 or a Fe chloride?
[Edited on 10-29-2018 by RogueRose]
[Edited on 10-29-2018 by RogueRose]
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j_sum1
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A few misconceptions here.
First, the ranking of the metals you are after is the "activity series" which is not precisely the same as reactivity.
It is best represented by the table of standard reduction potentials.
You will notice that lithium is further up the list than potassium even though potassium reacts more quickly. IOW, kinetics is not the same as the
driving force of a reaction.
The list I gave is not just a list of elements but rather a list of reactions: specifically reduction reactions under standard conditions. They are
ranked according to the electric potential of the electrons in the reaction. These potentials are listed.
In general, a metal high on the list will displace ions of a metal further down the list causing that species to precipitate out as a metal. It
matters little (in theory) what the anions are.
You will notice two bold reactions on the list: formation of H2 at -0.8277 and reduction of O2 at +0.401. Between these two are the reactions that can
take place in aqueous solution. These are the basis for your displacement reactions. So, for example, you are never going to form Al metal from a
solution. Nor is it sensible to use Na to displace Cu2+: it will react with the water not the copper ions.
I realise I have glossed oner some details. But I am typing on my phone so am being a bit brief. I hope this helps.
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Ubya
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j_sum1 is right, and if you want to apply "to the real world" those standars potentials (they are defined in standard condition 1Bar, 25°C, 1M) use
the Nernst equation, this way you can play a bit and let a reaction happen that wouldn't in standard conditions
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RogueRose
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Quote: Originally posted by j_sum1  | A few misconceptions here.
First, the ranking of the metals you are after is the "activity series" which is not precisely the same as reactivity.
It is best represented by the table of standard reduction potentials.
You will notice that lithium is further up the list than potassium even though potassium reacts more quickly. IOW, kinetics is not the same as the
driving force of a reaction.
The list I gave is not just a list of elements but rather a list of reactions: specifically reduction reactions under standard conditions. They are
ranked according to the electric potential of the electrons in the reaction. These potentials are listed.
In general, a metal high on the list will displace ions of a metal further down the list causing that species to precipitate out as a metal. It
matters little (in theory) what the anions are.
You will notice two bold reactions on the list: formation of H2 at -0.8277 and reduction of O2 at +0.401. Between these two are the reactions that can
take place in aqueous solution. These are the basis for your displacement reactions. So, for example, you are never going to form Al metal from a
solution. Nor is it sensible to use Na to displace Cu2+: it will react with the water not the copper ions.
I realise I have glossed oner some details. But I am typing on my phone so am being a bit brief. I hope this helps. |
Thanks, that is quite a reply on a phone! This has cleared some things up. That is quite a list to figure out.
So what would the numbers be (on the charge side) if I where to do this equation
CuSO4 + Fe -> FeSO4 + Cu
or
CuSO4 + Zn -> ZnSO4 + Zn
I'm pretty sure both of these reactions work but I don't see either on the list and not sure how decipher it. Do the positive voltages exchange with
a lower charge (does it have to be a negative or just lower positive charge).
What would be good to read to get a better understanding on how to make sense of the list, topic wise in chemistry.
Thanks again!
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j_sum1
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SO4 is just spectator. Ignore it.
Your first reaction:
Cu2+ + Fe --> Cu + Fe2+
Break it into two reactions:
Reduction
Cu2+ + 2e- --> Cu
Look up the reduction potential.
Oxidation
Fe --> Fe2+ + 2e-
You won't find this on the list. But you will find the reverse. So, switch the sign of the reduction potential to get the oxidation potential.
Add these two numbers together. If the result is negative you will have a reaction that occurs spontaneously.
I'll leave the Zn as an exercise for the reader.
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AJKOER
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A comment to quote from a source (Biochimica et Biophysica Acta (BBA) - Bioenergetics, Volume 1604, Issue 1, 18 April 2003, Pages 13-21, 'pH-dependent
redox potential: how to use it correctly in the activation energy analysis', at https://www.sciencedirect.com/science/article/pii/S000527280... upon scrolling down) relating to pH dependent redox potential, to quote:
" DXH <--> DX + H+ + e- (ll)
The designation D is used here for the redox center, for example, metal ion, and X for its protonable ligand, but such a division is in principle not
obligatory (see, e.g., quinol–semiquinone couple). Redox potential for reaction (II) is pH dependent"
where, as of yet, pH has not be mentioned. Another source, see 'pH Dependent Redox Couple: An Illustration of the Nernst Equation', by Mary M.
Walczak, et al, in J. Chem. Educ., 1997, 74 (10), p 1195, DOI: 10.1021/ed074p1195, link: https://pubs.acs.org/doi/10.1021/ed074p1195 and discussion at https://www.quora.com/How-does-pH-affect-reduction-potential .
Interestingly, in the first source, there is also some discussion depending on whether the reagent (in the current context, a metal) is not dissolved.
I do recall once seeing a very short table of standard reduction potentials with a column for acidic and basic values presented in an online ebook.
Here is an example of an article showing pH dependence, see https://febs.onlinelibrary.wiley.com/doi/pdf/10.1046/j.1432-... and another disclosing even a double pH dependent for thiol-containing
biomolecules (https://www.nature.com/articles/srep37596 ).
[Edited on 29-10-2018 by AJKOER]
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