# Electrode Potentials and Complexing

1. Dec 20, 2005

### Selectron

The standard electrode potential of a Cu/Cu^2+ half-cell is given as 0.34V. However, the half-cell will only have this electrode potential relative to the hydrogen electrode when the concentration of Cu^2+ ions is 1mol.dm^-3.

But won't most of the aqueous Cu^2+ ions form complexes, effectively lowering the concentration of Cu^2+ in the solution and so lowering the electrode potential of the half-cell (as. with fewer free Cu^2+ ions, more Cu atoms will be oxidised at the electrode, making the electrode potential of the half-cell more negative)?

So when the 0.34V figure is given, is this referring to a half-cell where the concentration of Cu^2+ ions *plus* the concentration of Cu complexes is 1mol.dm^-3 or where the concentration of the Cu^2+ ions *only* is 1mol.dm^-3?

P.S. Is it possible to make text superscripted in these forums, and if so, how do you do it?

2. Dec 20, 2005

### iansmith

Staff Emeritus
You can use latex for equations, superscript and subscripts
or you can use {sub}your text{/sub} for subscript and {sup}your text{/sup} for superscript. Just replace the {} with []

For example: writesubscript and writesuperscript

3. Dec 24, 2005

### GCT

although some complexes may form, such as with any charged ion in an aqueous solution, this won't significantly affect the direction through which the major reaction proceeds, think la chatelier's principle, it's sort of like the reaction between weak bases/acids with a strong acid/base, as long as you have the strong acid/base there all of the weak base/acid will be consumed in a stoichiometric titration.

4. Dec 25, 2005

### Staff: Mentor

For standard potential you need activity of Cu2+ to be exactly 1. If there are complexing agents they can change VERY SIGNIFICANTLY potential of the electrode - for example

E(Cu2+/Cu) = +0.34

but

E(Cu(NH3)42+/Cu) = -0.05.

Note that Cu2+/Cu means in fact that there is a copper (II) complex with water!

5. Dec 25, 2005

### Staff: Mentor

You are completely wrong. Adding complexing agents you may reverse the reaction direction.

6. Dec 25, 2005

### GCT

again borek, make sure you fully understand the context of the question and my reply, if you're having trouble understanding the english, read over the question/answer twice. I never said anything about adding complexing reagents such as in relation to EDTA titrations.

yes, it's called the nernst equation, in most experimental cases you simply won't have standard conditions especially during the course where the reaction proceeds. You make no significant point here.

Water forms an ordered structure around charged ions, it doesn't always act as a chelating agent or even a strong ligand for that matter, although it contributes to the hydrated radii. The question pertained specifically to whether an aqueous solution of copper in which the copper has an ordered structure with water would have a significant equilibrium which would contribute some of the copper not taking part in electrochemical reactions.

7. Dec 25, 2005

### GCT

that is you don't need to change the standard reduction potential values

what on earth is this supposed to mean?

8. Dec 26, 2005

### Selectron

OK thanks guys. I think I understand now.

9. Dec 26, 2005

### Staff: Mentor

Selectron never stated what other complexing agents can be present in the solution. He referred to "aqueous ions" and "concentration of Cu^2+ ions *plus* the concentration of Cu complexes". Thats very vague and wide description, fulfilled by all aqueous solutions containing all possible Cu2+ complexes.

You may be right in your reading of the question, as on the second reading it is ambiguous. Still your answer is at least partially wrong. First, all copper ions are complexed by water (not "some complexes may form" as you suggested), second, complexing can change direction of the reaction.

10. Dec 26, 2005

### GCT

no, experiments such as these are done all the time in elementary labs, and especially when you have a reference S.H.E. electrode, such complications don't exist.

no, complexing will not change anything unless you actually have a blocking agent of some sort which has a very strong affinity for the copper ions. There are no such problems with "complexing" in the real labs and they are done all the time without complications. The nernst equation is used and so are the standard potentials.

11. Dec 26, 2005

### Staff: Mentor

Complexing can change everything.

I don't have time atm to look for better example, but let's try with solution with [Cu2+] = 1M and [Ag+] = 0.01M. Put silver and copper wire into solution (don't connect them).

Potentials - as given by Nernst equation - are:

for Cu - E = 0.35 + 0.059/2 log([Cu2+ ]) = 0.35
for Ag - E = 0.80 + 0.059 log([Ag+ ]) = 0.68

As expected, silver deposits on the copper wire, reduced by copper.
Now, let's add enough thiosulfate to the solution to complex both copper and silver, and to obtain concentration of free thiosulfate of 1M. What happens to potentials?

Logs of overall stability constants are (L stands for S2O3):
for Cu(L)22- - 12.29
(no constant given in my tables for CuL, but it will not change a thing, as it probably means it is much smaller)

for Ag there are at least three complexes wih overall stability constants:
Ag(L)- - 8.82
Ag(L)23- - 13.67
Ag(L)35- - 14.2

To use Nernst equation (using standard potentials) we need to calculate concentration of free ions (or rather these complexed by water). They are given by the equation:

$$[Me] = \frac{C_{Me}}{1 + \beta_1 [L] + \beta_2 [L]^2 + \beta_3 [L]^3 + ...}$$

where [Me] is concentration of free ions, [L] is concentration of ligand, CMe is total concentration of metal - 1M and 0.01M respectively, $$\beta_x$$ are overall stability constants.

Putting known values into above we get

[Cu2+] = 5.1*10-13
[Ag+] = 4.9*10-17

and potentials - as given by Nernst equation:

for Cu - E = 0.35 + 0.059/2 log([Cu2+ ]) = -0.013
for Ag - E = 0.80 + 0.059 log([Ag+ ]) = -0.16

Look at these numbers - situation is reversed and we should expect copper to deposit on silver wire.

That's example of reversing the reaction direction with complexation, using common reagents, present in almost any lab.

Last edited: Dec 26, 2005
12. Dec 26, 2005

### GCT

don't connect them??! Why? Also you're deliberately adding a complexing agent, AGAIN, as I said, some complexing agents can act as blocking agents. But this has nothing to do with the topic....try to stay focused here borek. Also there are some problems even with you're current proposal, although it's not relevant to the original topic, I'll proceed to mention them later.

13. Dec 27, 2005

### Staff: Mentor

To have two different redox systems. If you connect the wires you will end with one system and both wires will be forced to have the same potential. That's not what we need.

"You are wrong and I will show it later" is a phrase I am getting used to. It is known method of making an impressions on the readers that you are right and your word was final. It assumes readers will not remember you have never answered. Like here: http://www.chemicalforums.com/index.php?board=8;action=display;threadid=5835. Unfortunately, I remember.

So, please show NOW, what is wrong with my approach. I am ready to learn something new and to admit I was mistaken if you PROVE it.

14. Dec 27, 2005

### GCT

Borek, you're in your own world aren't you? What the **** are you talking about? You have deliberately added a chelating/ligand to the situation, the original problem refers to copper in an aqueous solution, it doesn't get any simpler then that; and no the op does not need to change the standard redox potentials, just use the nernst equation. This has been established. What are you really getting at here? Prove WHAT? Why are you setting up all of these useless calculations? If you're trying to assert that certain complexing agents do decrease the maximum available power that a cell can put up, I can do the same, look into a standard analytical chemical text. Yes, add EDTA to the solution and you'll have decreased activity of the metal, you'll need a greater overpotential.

Better yet, simply mess around with the initial concentrations, I can reverse the direction of the reaction that way. You don't need any thiosulfate. All you're doing here is messing around with the concentrations, that's essentially what the Nernst equation implies, you can affect the spontaneity of the reaction by altering the concentrations. You don't even need to add a chelating agent, add more or less metal in varying compositions. Got it?

If you want to argue about something else, I'll be happy to have a peaceful discussion with you. Start a new thread in the chemistry subforum, the original query has been resolved.

15. Dec 27, 2005

### Bystander

No.

Yes.

Welcome to the wonderful world of "standard states." Hypothetical constructs used as references for thermodynamic measurements.

16. Dec 27, 2005

### Staff: Mentor

You have lost me. Is it finally possible to change reaction direction with complexation, or not?

And what does "copper in aqueous solution" stand for - is it by definition copper salt in pure water, or is it copper salt in water containing any other substances? I see ambiguity here.

17. Dec 27, 2005

### GCT

borek, please read the original question. I have said several times on several of my post here that, yes, if you deliberately add a chelating agent to the solution the required overpotential will be greater as well as the polarization potential. But this makes no difference on which standard potential to use with the nernst equation. So yes, you can change the direction of the reaction, you can affect the spontaneity of the reaction, by adding a blocking agent or chelating agent, or better yet simply add the two metals in varying composition. Adding more or less of the other can change the direction of the reaction as indicated by the nernst equation....

Copper in aqueous solution is simply the copper salt in water, with an innert electrolyte conjugate, denoted by $$Cu^{2+}_{aq}$$

please read my posts twice and understand them fully before you rant about it.