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According to my textbook, element $\ce{X}$ can oxidize element $\ce{Y}$ if $\ce{X}$ was lower in the activity series. Lead is lower than Tin in the activity series table in the textbook, but the following question contradict that rule and I can't figure out why:

Will the reaction $\ce{Sn(s) + Pb(NO3)2(aq) -> Sn(NO3)2(aq) + Pb(s)}$ occur?

The book reveals at the end that the answer is no, but doesn't explain why.

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up vote 4 down vote accepted

Are you certain that your statement from the textbook is correct? The cation of X will oxidize element Y if element Y is more active than X, which is usually labeled (may be towards the top or the bottom of the list). For reference, see the activity series at Wikipedia, where the more active metals are above the less active metals in the list.

How to use the activity series:

For example, what happens when iron metal is placed into an aqueous solution of copper (II) nitrate?

Let's write the potential equation:

$$\ce{2Fe(s) + 3Cu^{2+} -> 2Fe^{3+} + 3Cu(s)}$$

Since iron is more active than copper, iron will reduce copper cations (and copper cations will oxidize iron metal). The iron will dissolve.

What happens if copper metal is placed into an aqueous solution of iron (III) nitrate?

Since copper is less active than ion, copper metal will not reduce iron cations. No reaction occurs and the copper will not dissolve.

Your question

Lead is less active than tin in all activity series that I have seen. Lead metal should reduce tin cations. This reaction should work. It is possible that the textbook/solutions manual has a typo. I find them in my books all the time.

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Interestingly, I have a link for a video of that exact reaction. – Nicolau Saker Neto Jun 25 '14 at 1:09

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