I was just going through my last year's school textbooks and while I was reading this chapter on Redox, this above question crossed my mind. Shouldn't it be true that since ALL metals lose electrons to get stable, they should also get oxidized. And since they getting oxidized themselves, they should act as reducing agents... Always? If the answer is no, please provide an example of the exception.
$\begingroup$
$\endgroup$
0
Add a comment
|
2 Answers
$\begingroup$
$\endgroup$
Your reasoning is sound. All (pure) metals can act as reducing agents, but they do not have to all the time. There are a few examples of metals in negative oxidation states, where they would act as oxidising agent.
Probably the most prominent examples are Zintl phases, where alkaline or earth alkaline metals reduce post transition metals (or metalloids like antimony). One of the first reactions observed was the sodium lead reactions in liquid ammonia in 1907. Back in the day they proposed the following reaction equation: $$\ce{Na + excess~Pb ~->[\ce{NH3}][\mathrm{-50^\circ C}]~ Na+ + $“$~Pb2^{-}~$”$}$$ It was later found that the actual composition is $\ce{Na4Pb9}$. With crown ethers and/or cryptants it is possible to isolate discrete anions, too.
There are more examples like this, and a nice overview is given in John E. Ellis, Inorg. Chem., 2006, 45 (8), 3167–3186.
answered Apr 15, 2015 at 5:57
$\begingroup$
$\endgroup$
2
No, not always.
For example, in the classic Galvanic cell: 
In the above image, you have zinc metal acting as a reducing agent at the anode, true to your line of thinking. However, it is also the case that the copper ions at the cathode are being reduced.
This can happen quite frequently in electrochemistry as the reduction reaction that occurs is largely governed by the reduction potential of that reaction. In the Galvanic cell example, the copper reduction reaction has a higher standard reduction potential (+0.34V) compared to that of the zinc reduction (-0.76V), meaning that the copper ions are more likely to get reduced, and so the copper reduction will proceed over that of the zinc.
Image: http://chemwiki.ucdavis.edu/@api/deki/files/119/ZnCu_detail.gif?revision=1
-
2$\begingroup$ Elemental Cu is not being reduced. Cu ions are being reduced to elemental Cu. Cu 2+ ions are an oxidising agent as they can be reduced to Cu metal. Elemental Cu is a reducing agent as it can be oxidised to Cu ions. Under normal circumstances, elemental Cu cannot act as an oxidising agent. There are some cases where a very strong reducing agent can cause elemental metals to act as oxidising agents, forming anionic species. en.wikipedia.org/wiki/Zintl_phase gives for example 4(Na)+ . (Si4)4- If you don't consider Si a metal, there are similar compounds with Sn and Pb (and In and Tl) $\endgroup$ Commented Apr 15, 2015 at 10:36
-
$\begingroup$ You are right. I was being overly broad by not distinguishing between the two forms of copper. I'll change my answer to show that. $\endgroup$– ShafterCommented Apr 18, 2015 at 21:14