Is there a reason why this question's answer is treating it as if it is at standard conditions? Isn't it using 0.2M aqueous solutions, so the electrochemical series won't be valid for this question? The answer seems to ignore this.



While the electrochemical series does assume standard-state conditions, we will show that in certain cases it will also work for non-standard-state conditions. One recalls that the Nernst equation is given by $$E = E^\circ - \frac{RT}{nF}\,\ln Q,$$ and can be used to determine electrochemical cell potentials when the cells are not at standard state. This works directly for the second reaction you list, as $Q = \ce{[Sn^2+]}/\ce{[Cu^2+]} = 1$ and $\ln Q = 0$, so the second term in the Nernst equation vanishes, but will not work for the first reaction, for which $Q = \ce{[Sn^2+]}/\ce{[Ag+]^2} \neq 1$.

In this case, it is reasonable to assume that one knows both (i) that the correction due to the second term is relatively small, and (ii) that $\ce{Ag}$ and $\ce{Cu}$ are relatively inert metals, and strongly favor being in their elemental form relative to other metals. The quoted result then follows.

  • $\begingroup$ Thanks. Is this expected knowledge for a final year high school knowledge? Seems rather complex. $\endgroup$
    – Destudent
    Apr 9 '17 at 10:46
  • $\begingroup$ Yes---at least this was my experience in adv. chemistry class in high school. Your mileage may vary. $\endgroup$ Apr 9 '17 at 14:50

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