Electrochemical Series: Definition and uses | Electrochemistry Class 12

The Series obtained when the standard electrode potentials of different electrodes or semi-reactions occurring on them are in increasing order is called the electrochemical Series. All metals in this Series are located in decreasing order of their reactivity.

Main uses of electrochemical Series

1) The reactivity of metals can be compared with the help of electrical chemical grade.

In the Electrochemical Series which is located above the metal –

• Their electrode potentials are lower,
• Their reduction potentials are lower,
• Their oxidation potentials are high.
• The semi-reactions occurring on them have a higher tendency to oxidize.
• Those metals are more prone to oxidation.
• The tendency of those metals to give up electrons is high.
• These metals are more functional.

Therefore, the reactivity of metals decreases in increasing order of their electrode potentials and metals in the electrochemical series are placed in decreasing order of their reactivity.

Example: Fe is a more reactive metal than Cu.

Putting a more active metal in a solution of the cations of a less active metal frees the less active metal.

Example: Fe is located above Cu in the Electrochemical Series and is more reactive than Cu.

2Fe + 3CuSO₄ → Fe₂(SO₄)₃ + 3Cu

Therefore, by releasing Fe into the CuSO₄ solution, it liberates Cu.

2) With the help of the electrochemical Series, information is available about the ability of different metals to displace hydrogen from acids.

In the electrochemical Series, metal is located above hydrogen, compared to hydrogen.

• Their electrode potentials are lower,
• Their reduction potentials are lower,
• They have higher oxidation potential.
• They are prone to oxidation in semi-reactions.
• They have a higher tendency to oxidize.
• Those metals have a higher tendency to make cations.
• Those metals can displace hydrogen from acids.

Therefore, in the electrochemical category, the metal which is located above hydrogen, displaces hydrogen from the acids.

In other words, the metals whose electron potentials are lower in the electrode potential of hydrogen displace hydrogen from the metal acids. Metals whose electrode potentials are greater than the electrode potential of hydrogen do not displace hydrogen from the metal acids.

2Al + 6HCl → 2AlCl₃ + 3H₂

Zn + 2HCl → ZnCl₂ + H₂

Example: Aluminum and zinc lie above hydrogen in the electrochemical series. The standard electrode potentials of Al, Zn and hydrogen are 1.66, -0.76 and 0 volts respectively. Hence, Al and Zn displace hydrogen from HCl. The reaction speed of Al with HCl is faster than the reaction of Zn with HCl.

Copper is located below the hydrogen in the electrochemical series. The standard electrode potential of Cu is +0.34 volts which is higher than the standard electrode potential of hydrogen. Thus Cu cannot displace hydrogen from HCl.

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3) With the help of the electrochemical series, the oxidizing capacities of different oxidants and the oxidizing capacities of different catabolists can be compared.

From the electrode potential of a semi-reaction, it is known how much it is prone to oxidation or reduction. If the electrode potential of a semi-reaction is high then –

• The reduction potential of that semi-reaction will be higher
• In that semi-reaction, the tendency of reduction will be more
• The oxidizer used in that semi-reaction will be strong oxidizer.
• The reducing agent used in that half-reaction will be a weak reducing agent.

Similarly, if the electrode potential of a semi-reaction is low, the reducing agent used in that semi-reaction will be a strong oxidizer and the oxidizer used in that semi-reaction will be a weak oxidizer.

Example:

Zn²⁺ + 2e^– ⇌ Zn (E₀= -0.76Volt)

Cu²⁺ + 2e^– ⇌ Cu (E₀= +0.34Vold)

The above semi-reactions are written for reduction. Hence the voltage E₀ written with them shows their standard reduction potential. In the above semi-reactions, Zn²⁺ is Cu²⁺ oxidant and Zn is Cu oxidant.

Since the reduction potential of the first half-reaction (Zn²⁺ + 2e ⇌ Zn) is low, so in this half-reaction, oxidizer will be strong(Zn → Zn²⁺ + 2e), Zn is a strong oxidizer and Zn²⁺ is a weak oxidizer. Since the second half-reaction(Cu²⁺ + 2e ⇌ Cu) has a higher reduction potential, so this half-reaction(Cu²⁺ + 2e → Cu) will have a reduction, Cu²⁺ is a strong oxidizer and Cu is a weak oxidizer.

Following is the mixed equation of both the above semi-reactions.

Zn + Cu²⁺ ⇌ Zn²⁺ + Cu

Since Zn is a fast reducing agent and Cu²⁺ is a strong oxidizer, this reaction is possible in the right direction. This reaction is not possible in the left direction because Zn²⁺ is a weak oxidizer and Cu is a weak oxidizer.

Zn + Cu²⁺ → Zn²⁺ + Cu

When Zn rods are added to CuSO₄ solution, Zn will dissolve in solution and Cu will precipitate but there will be no action if Cu rod is added to ZnZO₄ solution.

4) The E.M.F. of a galvanic cell can be calculated from the electrochemical Series.

E.M.F. of a galvanic cell = potential of positive electrode – potential of negative electrode

Standard electrode potentials of various electrodes are given in the electrochemical series. Therefore, with its help the E.M.F. of a galvanic cell can be calculated.

Example: If a galvanic cell has zinc and copper rods immersed in one molar solution of ZnSO₄ and CuSO₄, then the zinc rod has negative charge and negative potential and the copper rod has positive charge and money potential. The device consists of a Cu positive electrode and a Zn negative electrode. Cu has E₀ = +0.34 volts and Zn has E₀ = -0.76 volts.

Therefore

E_Cell = E_Cathode + E_Anode

= E_Cu – E_Zn 

          = + 0.34 – (-0.76)

= 1.10 Volt 

Question: What is the reason that copper immersed in copper sulphate solution freezes copper?

Answer:

First Method: In electrochemical series iron is stable above copper. Hence iron is a more functional metal than copper. The more reactive metal displaces the less reactive metals from the solutions of their salts. Therefore, the iron reaction will be followed by the reaction of copper sulphate in the solution.

Fe + CuSO₄ → FeSO₄ + Cu

As copper becomes free, it gets deposited on iron nails.

Second Method: In the electrochemical series, iron is located above copper. Hence the electrode potential of iron is less than the electrode potential of copper.

Fe²⁺ + 2e ⇌ Fe

Cu²⁺ + 2e ⇌ Cu

The reduction potential of the first half-reaction is low and the oxidation potential is high. The reduction potential of the second half-reaction is high. And oxidation potential is low. Hence Fe will be oxidized to Fe and Cu will be reduced to Cu.

Fe → Fe²⁺ + 2e

Cu²⁺ + 2e → Cu

Copper sulphate solution contains Cu copper cations. Adding iron to it will cause oxidation of iron and reduction of Cu.

Fe + Cu²⁺ → Fe²⁺ + Cu

Fe + CuSO₄ → FeSO₄ + Cu

As copper(Cu) becomes free, it gets deposited on iron nails.