Activity Series of Metals

The activity series of metals ranks metals by their tendency to lose electrons and undergo oxidation. This tendency is directly related to their standard electrode potentials (E°). Metals higher in the series are more reactive, more easily oxidized, and thus have more negative standard reduction potentials.

Relation to Electrode Potentials

In the electrochemical series (a list of half-reactions with their standard electrode potentials), metals that are strong reducing agents appear at the top of the activity series with large negative \(E^\circ\) values. Conversely, metals with high (positive) \(E^\circ\) values are less reactive and found lower in the activity series.

Example segment from the electrochemical series:

\( \text{K}^+ + e^- \rightarrow \text{K (s)} \qquad E^\circ = -2.93 \, \text{V} \)
\( \text{Zn}^{2+} + 2e^- \rightarrow \text{Zn (s)} \qquad E^\circ = -0.76 \, \text{V} \)
\( \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu (s)} \qquad E^\circ = +0.34 \, \text{V} \)
\( \text{Ag}^+ + e^- \rightarrow \text{Ag (s)} \qquad E^\circ = +0.80 \, \text{V} \)

From this, we observe that potassium is a much stronger reducing agent than silver, which is why it appears higher in the activity series.

Applications

Displacement reactions: A metal will displace another from its salt solution if its standard electrode potential is more negative.
Redox spontaneity: Comparing two half-cell potentials tells us whether a redox reaction is spontaneous under standard conditions.
Predicting cell reactions: Activity series and electrode potentials allow you to predict the direction of electron flow in galvanic cells.

Consider the reaction:

\[ \text{Zn (s)} + 2\text{Ag}^+ (aq) \rightarrow \text{Zn}^{2+} (aq) + 2\text{Ag (s)} \]

Solution:

\( E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76 \, \text{V} \)
\( E^\circ_{\text{Ag}^+/\text{Ag}} = +0.80 \, \text{V} \)
\( E^\circ_{\text{cell}} = 0.80 - (-0.76) = +1.56 \, \text{V} \)
The positive cell potential confirms the reaction is spontaneous. Zinc displaces silver.

\[ \text{Cu (s)} + \text{Zn}^{2+} (aq) \rightarrow \text{Cu}^{2+} (aq) + \text{Zn (s)} \; ? \]

Solution:

\( E^\circ_{\text{Cu}^{2+}/\text{Cu}} = +0.34 \, \text{V} \)
\( E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76 \, \text{V} \)
\( E^\circ_{\text{cell}} = -0.76 - 0.34 = -1.10 \, \text{V} \)
The negative potential shows the reaction is not spontaneous. Copper cannot displace zinc.

Thus, the activity series and electrochemical series are two sides of the same coin — one qualitative, the other quantitative — both essential for mastering redox chemistry and electrochemistry.

Written by Thenura Wickramaratna