Place a strip of metallic zinc in an aqueous solution of copper nitrate as shown in Figure, for about one hour. You may notice that the strip becomes coated with reddish metallic copper and the blue colour of the solution disappears.
Formation of Zn2+ ions among the products can easily be judged when the blue colour of the solution due to Cu2+ has disappeared.
If hydrogen sulphide gas is passed through the colourless solution containing Zn2+ ions, appearance of white zinc sulphide, ZnS can be seen on making the solution alkaline with ammonia.
The reaction between metallic zinc and the aqueous solution of copper nitrate is:
Zn(s) + Cu2+ (aq) → Zn2+ (aq) + Cu(s) -------(4)
In above reaction, zinc has lost electrons to form Zn2+ and, therefore, zinc is oxidised.
Evidently, now if zinc is oxidised, releasing electrons, something must be reduced, accepting the electrons lost by zinc. Copper ion is reduced by gaining electrons from the zinc.
Reaction may be rewritten as:
At this stage we may investigate the state of equilibrium for the reaction represented by equation (4). For this purpose, let us place a strip of metallic copper in a zinc sulphate solution.
No visible reaction is noticed and attempt to detect the presence of Cu2+ ions by passing H2S gas through the solution to produce the black colour of cupric sulphide, CuS, does not succeed.
Cupric sulphide has such a low solubility that this is an extremely sensitive test; yet the amount of Cu2+ formed cannot be detected.
We thus conclude that the state of equilibrium for the reaction (4) greatly favours the products over the reactants. Let us extend electron transfer reaction now to copper metal and silver nitrate solution in water and arrange a set-up as shown in Figure.
The solution develops blue colour due to the formation of Cu2+ ions on account of the reaction:
Here, Cu(s) is oxidised to Cu2+(aq) and Ag+(aq) is reduced to Ag(s). Equilibrium greatly favours the products Cu2+ (aq) and Ag(s).
By way of contrast, let us also compare the reaction of metallic cobalt placed in nickel sulphate solution. The reaction that occurs here is:
At equilibrium, chemical tests reveal that both Ni2+(aq) and Co2+(aq) are present at moderate concentrations. In this case, neither the reactants [Co(s) and Ni2+(aq)] nor the products [Co2+(aq) and Ni (s)] are greatly favoured.
This competition for release of electrons incidently reminds us of the competition for release of protons among acids. The similarity suggests that we might develop a table in which metals and their ions are listed on the basis of their tendency to release electrons just as we do in the case of acids to indicate the strength of the acids.
As a matter of fact we have already made certain comparisons. By comparison we have come to know that zinc releases electrons to copper and copper releases electrons to silver and, therefore, the electron releasing tendency of the metals is in the order: Zn>Cu>Ag.
We would love to make our list more vast and design a metal activity series or electrochemical series. The competition for electrons between various metals helps us to design a class of cells, named as Galvanic cells in which the chemical reactions become the source of electrical energy.