2.5 The Activity Series

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As we mentioned earlier (Section 2.3), elements differ in their reactivity. Some elements are so reactive that they do not occur naturally in their elemental form; rather they only exist in nature as part of a compound. Other elements are essentially completely unreactive and are chemically inert (the Noble gases).

An activity series is often useful is helping to predict whether certain reactions will occur. An activity series list metals and various other elements in order of their reactivity, with the most reactive elements at the top and the least reactive of the series at the bottom. For any two metals, the metal listed higher in the table is the most readily oxidized.

Activity Series of Metals in Aqueous Solutions
most active (easily oxidized – readily lose electrons))
lithium	\(\ce{Li}\)	These metals displace hydrogen from water \(\ce{Ca(s) + 2H2O(l) -> Ca(OH)2 + H2(g)}\) These elements are very reactive and react readily to form compounds
potassium	\(\ce{K}\)
barium	\(\ce{Ba}\)
calcium	\(\ce{Ca}\)
sodium	\(\ce{Na}\)
		These metals displace hydrogen from acids \(\ce{Zn(s) + HCl(aq) -> ZnCl2 + H2(g)}\)
magnesium	\(\ce{Mg}\)
aluminum	\(\ce{Al}\)
zinc	\(\ce{Zn}\)
chromium	\(\ce{Cr}\)
iron	\(\ce{Fe}\)
cadmium	\(\ce{Cd}\)
nickel	\(\ce{Ni}\)
tin	\(\ce{Sn}\)
lead	\(\ce{Pb}\)

hydrogen	\(\ce{H}\)

copper	\(\ce{Cu}\)	These metals do not displace hydrogen from acids or water These elements are more stable, and form compounds less readily than do those higher in the table.
silver	\(\ce{Ag}\)
mercury	\(\ce{Hg}\)
platinum	\(\ce{Pt}\)
gold	\(\ce{Au}\)
least active

Some examples of how reactions can be predicted using an activity series:

Any metal on the list can be oxidized by the ions of elements below it in the chart. For example copper is above silver in the chart. Thus, copper metal will be oxidized by silver ions as in the following example:

\(\ce{Cu(s) + 2AgNO3(aq) -> Cu(NO3)2(aq) + 2Ag(s)}\)

Would it be possible to store a silver spoon in a zinc nitrate solution? That is, will the following reaction occur:

\(\ce{2Ag(s) + Zn(NO3)2(aq) -> 2AgNO3(aq) + Zn(s)}\)

Since silver is below zinc on the chart, silver metal will not be oxidized by zinc. Therefore it would be safe to store the silver spoon in the zinc solution; the silver spoon will not undergo oxidation and corrode.

Would it be possible to store a silver nitrate solution in a copper container? That is, will the following reaction occur:

\(\ce{Cu(s) + 2AgNO3(aq) -> Cu(NO3)2(aq) + 2Ag(s)}\)

We do not want this reaction to occur, since the copper container would corrode.

From the table we find that Cu is a more easily oxidized than is Ag. In other words, copper will undergo oxidation and the reaction will occur. Therefore, it is not recommended to store AgNO₃in a copper container.