Question (HELP NEEDED): Electrolytic Cells

[pm123]

Hi,
I had an urgent question regarding electrolytic cells.
In class, we observed an electrolytic cell, where the cathode was copper, anode was nickel and the electrolyte was IM Nickel sulfate solution.
We had passed 1 amp current for 10 minutes, and nickel had coated on the copper electrode just like we predicted.
But for some reason, the nickel electrode did not corrode. Why is this the case? In the textbook, it says that the concentration of the electrolyte does not change if it is the same as the electrode (i.e. Nickel sulfate solution and nickel electrode), as the nickel electrode corrodes and hence supplies the solution with the nickel ions. Or have I understood this wrong?
Also, what would happen if after electroplating for several minutes, the connections of the wires were reversed? Copper electrode would corrode as it becomes anode and oxidises, but would nickel coat on the nickel metal electrode or copper metal? And why is this the case?
Thanks for the help!!

[Mattjbr2]

Here you go, friend http://media.pearsoncmg.com/bc/bc_0media_chem/chem_sim/html5/Electro/Electro.php
Run each of those simulations and see what happens! If you have any more questions let me know

[pm123]

Here you go, friend http://media.pearsoncmg.com/bc/bc_0media_chem/chem_sim/html5/Electro/Electro.php
Run each of those simulations and see what happens! If you have any more questions let me know

Thank you! That really did help.
The simulation did have the nickel electrode corroding away, however, the experiment did not show this as this was happening at a very small level - would I record this as an observation? Or would I not?
Also, I tried reversing the wire connections, but it was not letting me make copper the anode, nickel the cathode in a nickel ion electrolyte. What would the overall outcome be in this scenario? Would copper ions from anode or nickel ions in solutoin reduce at cathode?

[Mattjbr2]

Thank you! That really did help.
The simulation did have the nickel electrode corroding away, however, the experiment did not show this as this was happening at a very small level - would I record this as an observation? Or would I not?
Also, I tried reversing the wire connections, but it was not letting me make copper the anode, nickel the cathode in a nickel ion electrolyte. What would the overall outcome be in this scenario? Would copper ions from anode or nickel ions in solution reduce at cathode?

If the experiment didn't show corrosion then you note that and provide possible reasons as to why that was the case. You don't under any circumstances falsify observations or change your data in order to fit theoretical expected observations (as I've seen a lot of students do to make their results look good  ).
Let us think methodically about what occurs in any cell.
To know what happens in our cell, we note down every single substance available in the cell. In the Ni(s) cathode / Cu(s) anode electrolytic cell we have: Ni(s), Ni²⁺(aq), Cu(s), H₂O(l) and SO₄^2-(aq).
The half equations are:
SO₄^2-(aq) + 10H⁺ + 8e^- <--> H₂S(g) +4H₂O(l), +1.7V
O₂(g) + 4H⁺(aq) + 4e^- <-->  2H₂O(l), +1.23V
Cu²⁺(aq) + 2e^- <--> Cu(s), +0.34V  (+) Anode
Ni²⁺(aq) + 2e^- <--> Ni(s), -0.25V  (-) Cathode
Ni²⁺(aq) + 2e^- <--> Ni(s), -0.25V
2H₂O(l) + 2e^- <--> H₂(g) + 2OH^-(aq), -0.83V

The substances we have are in bold. I included sulfate (not on our electrochemical series) just for educational purposes.

Cathode: At the cathode we have a build up of negative charge due to the force of incoming electrons from the external power source. The strongest oxidant available at this site will attract these electrons the most and will reduce. What oxidants are available? We have negative sulfate anions, positive nickel cations and water. Sulfate immediately is attracted to the positive electrode (because sulfate is negatively charged) and the positive electrode is all the way on the other side of our cell, so sulfate is no longer at the site of reduction and so it doesn't reduce. It's attracted to the positive copper anode. We're left with nickel ions and water. Nickel ions have a larger standard reduction potential than water and so they will reduce to form nickel solid. The nickel cathode hence gets coated in nickel and gains mass. The greenish colour of the solution will be reduced as nickel ions are reduced.
Anode: Here we have a build up of positive charge due to the force of receding electrons. The strongest reductants will be oxidised. What reductants do we have? We have water<solid copper<solid nickel. The nickel is at the other end of the cell, so it won't be oxidised here. Solid copper has a higher Standard oxidation potential (opposite to what VCAA gives us in our data book, which are Standard reduction potentials). We're left with solid copper as the strongest reductant at the anode. So copper will oxidise to form copper ions. The anode will lose mass. Sulfate and copper ions don't form a precipitate and so the electrolyte will turn blue around the anode.

As the reaction progresses, the concentration of nickel ions in the electrolyte decreases and the concentration of copper ions increase. The electrolyte will have changed from greenish to blue. Eventually it'll come to a point where the copper anode has mostly corroded and the electrolyte is almost entirely composed of CuSO₄(aq). Once the anode has completely corroded, the circuit is broken and the external power source isn't able to force electrons into the nickel cathode any more. Now you effectively have a piece of nickel metal dipped into a copper(ii)sulfate solution. A spontaneous single displacement reaction will occur and nickel will be oxidised into the solution while copper will be reduced and plated onto the corroding nickel plate. Now the blue solution will turn greenish once more.

Don't take my word for it 100% though, as I'm not an authority on VCE chem. My teacher is. I'll email her and ask her if what I said has any inaccuracies, then I'll update this post with her response.

[pm123]

If the experiment didn't show corrosion then you note that and provide possible reasons as to why that was the case. You don't under any circumstances falsify observations or change your data in order to fit theoretical expected observations (as I've seen a lot of students do to make their results look good  ).
Let us think methodically about what occurs in any cell.
To know what happens in our cell, we note down every single substance available in the cell. In the Ni(s) cathode / Cu(s) anode electrolytic cell we have: Ni(s), Ni²⁺(aq), Cu(s), H₂O(l) and SO₄^2-(aq).
The half equations are:
SO₄^2-(aq) + 10H⁺ + 8e^- <--> H₂S(g) +4H₂O(l), +1.7V
O₂(g) + 4H⁺(aq) + 4e^- <-->  2H₂O(l), +1.23V
Cu²⁺(aq) + 2e^- <--> Cu(s), +0.34V  (+) Anode
Ni²⁺(aq) + 2e^- <--> Ni(s), -0.25V  (-) Cathode
Ni²⁺(aq) + 2e^- <--> Ni(s), -0.25V
2H₂O(l) + 2e^- <--> H₂(g) + 2OH^-(aq), -0.83V

The substances we have are in bold. I included sulfate (not on our electrochemical series) just for educational purposes.

Cathode: At the cathode we have a build up of negative charge due to the force of incoming electrons from the external power source. The strongest oxidant available at this site will attract these electrons the most and will reduce. What oxidants are available? We have negative sulfate anions, positive nickel cations and water. Sulfate immediately is attracted to the positive electrode (because sulfate is negatively charged) and the positive electrode is all the way on the other side of our cell, so sulfate is no longer at the site of reduction and so it doesn't reduce. It's attracted to the positive copper anode. We're left with nickel ions and water. Nickel ions have a larger standard reduction potential than water and so they will reduce to form nickel solid. The nickel cathode hence gets coated in nickel and gains mass. The greenish colour of the solution will be reduced as nickel ions are reduced.
Anode: Here we have a build up of positive charge due to the force of receding electrons. The strongest reductants will be oxidised. What reductants do we have? We have water<solid copper<solid nickel. The nickel is at the other end of the cell, so it won't be oxidised here. Solid copper has a higher Standard oxidation potential (opposite to what VCAA gives us in our data book, which are Standard reduction potentials). We're left with solid copper as the strongest reductant at the anode. So copper will oxidise to form copper ions. The anode will lose mass. Sulfate and copper ions don't form a precipitate and so the electrolyte will turn blue around the anode.

As the reaction progresses, the concentration of nickel ions in the electrolyte decreases and the concentration of copper ions increase. The electrolyte will have changed from greenish to blue. Eventually it'll come to a point where the copper anode has mostly corroded and the electrolyte is almost entirely composed of CuSO₄(aq). Once the anode has completely corroded, the circuit is broken and the external power source isn't able to force electrons into the nickel cathode any more. Now you effectively have a piece of nickel metal dipped into a copper(ii)sulfate solution. A spontaneous single displacement reaction will occur and nickel will be oxidised into the solution while copper will be reduced and plated onto the corroding nickel plate. Now the blue solution will turn greenish once more.

Don't take my word for it 100% though, as I'm not an authority on VCE chem. My teacher is. I'll email her and ask her if what I said has any inaccuracies, then I'll update this post with her response.

Thanks for the reply.
I am not too sure about the displacement reaction that was mentioned. One of the questions I have come across in the textbook suggests something along these lines:
The Copper electrode begins to corrode (anode) as it is oxidising into Cu2+(aq) ions. Thus, its concentration increases in the electrolyte. As it is a stronger oxidant than the Ni2+(aq) ions in the electrolyte, it instead reduces/precipitates onto the cathode rather than the nickel ions. I am not sure if that means that first nickel would be coated then copper, or if it will just be a coating of copper....