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I just finished a lab experiment in which I built a voltaic cell with copper (II) sulphate and zinc sulphate. My experiment was to identify the correlation between the concentration of copper (II) sulphate and the voltage generated.

  • I started with a 200 ml solution of 0.5 molar copper (II) sulphate and a 200 ml solution of 0.5 molar zinc sulphate solution (which is constant).
  • I measured the voltage of the cell at 0.5 M for both and I got about 1.088 V.
  • I added 1g of copper (II) sulphate anhydrous (crystal) to the copper (II) sulphate solution after every trial (which effectively increases the concentration of copper (II) sulphate (the cathode)) but I realized that each time I did that the voltage decreased. For context, I added 1g 18 times and I ended up with a voltage of 1.029 V.

So TLDR, increasing the concentration of copper (II) sulphate (the cathode) decreased the voltage of my voltaic cell instead of increasing it.

A few other students in my class did the same experiment (slight differences in methodology) and their voltage increased, which supports my original hypothesis of voltage increasing with concentration and also supports the Nernst equation (pretty sure).

Can anyone PLEASE help me understand why the voltage decreased when I increased the concentration of copper (II) sulphate?

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    $\begingroup$ There is possibility the activity coefficient of copper ions decreases faster than concentration increases, leading to decreasing of ion activity. Using concentrations instead of activities is just approximation, that is simple enough for school, but diverges from reality. // Did the students use the same/similar concentration range? Did the voltage with the given solution have any time trend? $\endgroup$
    – Poutnik
    Commented Nov 30, 2023 at 4:08
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    $\begingroup$ There are basically 2 options, implying you have all used the same or overlapping concentration range. 1/ it happened to all, so the activity thing is the reason. 2/ It happened just to you, so there is some error on your part we cannot remotely identify. // If there is possibility to repeat the experiment, use 10-50 times more diluted solution. 0.5 M is a lot. $\endgroup$
    – Poutnik
    Commented Nov 30, 2023 at 4:20
  • $\begingroup$ What sort of salt bridge have you used ? $\endgroup$
    – Maurice
    Commented Nov 30, 2023 at 9:57
  • $\begingroup$ Anhydrous copper sulfate is a white powder, not a crystal (as you described) $\endgroup$
    – Maurice
    Commented Nov 30, 2023 at 10:01
  • $\begingroup$ How do you explain that your students obtain a result which is contrary to yours ? $\endgroup$
    – Maurice
    Commented Nov 30, 2023 at 10:18

1 Answer 1

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Nernst eqn says that,

$$E_c=E^o_c-\frac{0.059}{n}logQ_c$$.

Now $Q_c = \frac{products}{reactants}$

If I take the cell reaction $\ce{Zn + Cu^2+ -> Zn^2+ + Cu}$,

$Q_c=\frac{[\ce{Zn^2+}](anode)}{[\ce{Cu^2+}](cathode)}$

So I must make sure that my $Q_c$ must be increased to lower the voltage. This can be made possible by increasing anode concentration and decreasing cathode concentration.

On the contrary, if I make cathode more concentrated and anode less concentrated, then $Q_c$ must be decreased, thus increasing EMF of the cell.

There is 1 possibility, that you interchanged both the cathode and anode and made the 1st scenario happen rather than the 2nd scenario which you were looking for.

SOURCE: NCERT

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  • $\begingroup$ Comments have been moved to chat; please do not continue the discussion here. Before posting a comment below this one, please review the purposes of comments. Comments that do not request clarification or suggest improvements usually belong as an answer, on Chemistry Meta, or in Chemistry Chat. Comments continuing discussion may be removed. $\endgroup$
    – Buck Thorn
    Commented Apr 23 at 18:49

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