Competitive Electron Transfer Reactions

We have seen that redox reactions involve electrons moving from one species to another. But here is a natural follow-up question: if two different metals are both capable of releasing electrons, which one releases them more readily? Can we rank metals by how eagerly they give up electrons? It turns out we can, and the way to do it is surprisingly hands-on: simply put one metal into a solution containing another metal’s ions and watch what happens.

The Zinc vs. Copper Experiment

Picture dropping a strip of metallic zinc into an aqueous copper nitrate solution, which has a distinct blue colour because of dissolved $Cu^{2+}$ ions. Leave it for about an hour and two striking changes appear:

• The zinc strip picks up a reddish coating of metallic copper.
• The blue colour of the solution fades and eventually disappears.

What is happening at the atomic level? Zinc atoms on the surface of the strip are releasing electrons and entering the solution as $Zn^{2+}$ ions. Those released electrons are snapped up by $Cu^{2+}$ ions floating in the solution, which get reduced to solid copper metal and plate themselves onto the strip. The overall reaction is:

$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s) \qquad \text{(7.15)}$

Fig 7.1: Redox reaction between zinc and aqueous copper nitrate solution, showing the progressive deposition of copper on the zinc strip

Zinc loses electrons (oxidation), and copper ions gain electrons (reduction). But how do we know $Zn^{2+}$ ions really formed? Once the blue colour has vanished (confirming $Cu^{2+}$ is gone), pass hydrogen sulphide ($H_2S$) gas through the now-colourless solution and make it alkaline with ammonia. A white precipitate of zinc sulphide ($ZnS$) appears, confirming that $Zn^{2+}$ ions are present. This is an extremely sensitive test, so even trace amounts of $Zn^{2+}$ would show up.

Where Does the Equilibrium Sit?

An important question is whether this reaction goes nearly to completion or just partway. To find out, try the reverse experiment: place a copper strip into zinc sulphate solution. Nothing visible happens. Even the ultra-sensitive $H_2S$ test for $Cu^{2+}$ ions comes back negative. Cupric sulphide ($CuS$) has such low solubility that it would show a black colour at the faintest trace of $Cu^{2+}$, yet none is detected.

The conclusion is clear: the equilibrium of reaction (7.15) overwhelmingly favours the products. Zinc pushes its electrons onto copper ions so effectively that the reverse process is virtually non-existent.

The Copper vs. Silver Experiment

Now try the next matchup. Place a copper strip into an aqueous silver nitrate solution. This time the solution develops a blue colour (a signature of $Cu^{2+}$ ions forming), and silver metal deposits on the copper strip. The reaction is:

$Cu(s) + 2Ag^+(aq) \rightarrow Cu^{2+}(aq) + 2Ag(s) \qquad \text{(7.16)}$

Fig 7.2: Redox reaction between copper and aqueous silver nitrate solution, showing the developing blue colour and silver deposition on the copper strip

Copper is oxidised (it loses electrons and enters the solution as $Cu^{2+}$), while silver ions are reduced (they gain electrons and come out of solution as solid silver). Once again, the equilibrium strongly favours the products: copper releases electrons to silver ions almost completely.

The Cobalt vs. Nickel Experiment: A Contrast

Not every displacement reaction goes nearly to completion. When metallic cobalt is placed in nickel sulphate solution, the reaction that occurs is:

$Co(s) + Ni^{2+}(aq) \rightarrow Co^{2+}(aq) + Ni(s) \qquad \text{(7.17)}$

Chemical tests at equilibrium reveal that both $Ni^{2+}$ and $Co^{2+}$ are present at moderate concentrations. Neither the reactant side nor the product side is strongly favoured. This tells us something important: cobalt and nickel have fairly similar tendencies to release electrons. Neither metal dominates the competition the way zinc dominates copper.

Building a Ranking: The Metal Activity Series

These three experiments reveal a clear pattern. Zinc pushes electrons onto copper ions almost completely, and copper pushes electrons onto silver ions almost completely. Chain these results together, and the order of electron-releasing tendency becomes:

$Zn > Cu > Ag$

Zinc gives up electrons most readily, copper comes next, and silver is the least willing to part with its electrons.

This idea of ranking metals by how easily they release electrons has a strong parallel in acid chemistry, where acids are ranked by how easily they release protons. Just as we build a table of acid strengths, we can build a table of metal “electron-releasing strengths.” This ranking is called the metal activity series (also known as the electrochemical series). It lists metals from those that release electrons most eagerly at the top to those that hold on tightest at the bottom.

From Beaker Experiments to Electrical Energy: Galvanic Cells

The competition between metals for electrons is not just an academic exercise. When two metals with different electron-releasing tendencies are connected through an external circuit, the spontaneous flow of electrons from the more active metal to the less active one creates an electric current. Devices that harness this principle are called Galvanic cells (named after Luigi Galvani). In a Galvanic cell, the chemical energy stored in the tendency difference between two metals is converted into electrical energy.

The zinc-copper pair we studied above, for example, is the basis of one of the most famous electrochemical cells. The details of how these cells are constructed and the quantitative relationships that govern them are explored in Class XII.