Electrochemical Series: Definitions, 12th Notes, Examples, Characteristics, Applications

What is Electrochemical Series?

An activity series, also known as an electrochemical series, is a collection of components arranged in an ascending order of increasing electrode potential values. The series was developed by comparing and measuring the Standard Hydrogen Electrode (SHE) in relation to the potential electrodes. All electrodes (metals and nonmetals) are grouped in an electrochemical series based on their distinct values of standard oxidation and reduction potential when in contact with their ions. The voltage measured when the half cell is connected to the standard hydrogen electrode under standard circumstances yields the standard electrode potential.

Features of the Electrochemical Series

The electrode potential of hydrogen (standard hydrogen potential, SHE) is 0.00. In the electrochemical series, half-cells (element/ion pairs) with a highly positive electrode potential are at the top. They are potent antioxidants. A reducing agent is a negative potential half-cell; the stronger the reducing power, the more negative the value. Metals are electropositive in general, whereas nonmetals are electronegative. 

The most active nonmetals are at the top of the column. As a result, the responsiveness is lowest along the center. The series of lower metals can diminish the upper metals. Similarly, nonmetals at the top of the series can oxidize metals, whereas nonmetals at the bottom of the series can oxidize metals.

Characteristics of the Electrochemical Series

Substances with higher reducing potentials than hydrogen are rated higher and have lower standard reduction potentials. Elements are organized in such a way that their strength as reducing agents decreases from top to bottom, while their strength as an oxidizing agents increases from bottom to top. All chemicals with positive reduction potential values that are positioned below hydrogen in the series are weaker reducing agents than hydrogen. Substances that are more powerful oxidizers than the H+ ion are listed after hydrogen in the series. Metals near the top (with large negative standard reduction potential values) quickly lose electrons. These are metals that are active. The nonmetals at the bottom (with high positive standard reduction potential values) readily receive electrons. These are nonmetals that are active. 

Standard Electrode Potential Value 

The value of the electrode potential is computed as follows-

1. A metal (or nonmetal) electrode and its ions are linked to the Standard Hydrogen Electrode (SHE). 
2. SHE is made of H2 and H+ ions. 
3. n typical settings, a characteristic voltage is measured across the electrodes depending on the metal and ions utilized. 
4. For the particular metal/ion combination, this is referred to as the "standard electrode potential value." 

Significance of Electrochemicals Series

1.   Cell EMF Calculation

Each electrochemical cell is made up of two half-cells, one for each electrode. Each half-cell passes through a reaction, one of which is oxidation and the other is reduction. Potentials, notably oxidation and reduction potentials, are associated with each process. Cell EMF (Ecell) is the sum of the oxidation and reduction potentials of the cell. It assesses the impulsiveness of the complete intracellular response. It is also a measure of how much work a cell can perform. The Electrochemical series aids in the measurement of a cell's EMF by obtaining readings of the half-standard cell's electrode potential and adding them correctly-

Eo Cell = Eo red + EoOX

where E∘red is the reducing half-standard cell's reduction potential and E∘oxd is the oxidizing half-standard cell's reduction potential.

2. Measuring Reaction Spontaneity

The electromotive force of the corresponding reaction is directly connected to the feasibility or spontaneity of the corresponding reaction. If the EMF of the cell is positive, the reaction is spontaneous. If the electromotive force of the cell is negative, the reaction is not spontaneous. By studying the reactants and products, one can assess if the redox reaction happens spontaneously. Create equations for both the reduction and oxidation halves of the process. Then, according to the electrochemical sequence, they add their typical electrode potentials. The EMF produced by the cell reveals if the reaction was spontaneous.

3. Calculating Gibbs Free Energy

Another indicator of reaction spontaneity is Gibbs free energy (ΔG∘cell). It is connected to the electromotive force (E∘cell) of the cell as follows-

ΔG∘cell = −nFE∘cell

where n denotes the number of electrons participating in the process and F denotes the Faraday constant, which is equal to 96.485 Coulomb-mol-1.

Based on the cell's EMF coding, if the EMF is negative, the Gibbs free energy is positive, and the reaction is not spontaneous. The reaction is spontaneous if the cell's EMF is positive and the Gibbs free energy is negative.

Electronegative and Electropositive Elements

Electropositive elements (other than hydrogen) are those that have a stronger tendency to lose electrons in solution. Elements that gain electrons are also known to be electronegative. They are commonly found in the series below the element hydrogen. In any event, we may deduce the sequence in which metals will replace one another by looking at the electrochemical series. As a result, electropositive metals often replace hydrogen from acids.

Electrochemical Series Chart

The Electrochemical Series chart is a quick way to see how similar and different metals are. Elements that are near each other tend to have similar qualities, but those that are far away have distinct dissimilarities. In other words, metals that are far apart in the electrochemical series will have a higher susceptibility to corrosion than metals that are close together.

In the most basic terms, metals that reside further apart on this scale will react with a greater tendency for corrosion than metals that are close together. (For example, zinc and copper are widely apart on the scale.) This implies that a copper pipe would never drain water onto a zinc-coated roof.

Read more about First 20 Elements of Periodic Table, Electrophile and Nucleophile, and Versatile Nature of Carbon

Application of the Electrochemical Series

Based on their location in the series relative to the hydrogen electrode, the electrochemical series may be used to predict certain characteristics of an element. The following are the applications of the electrochemical series-

1. Strengthening and Oxidizing

The electrochemical series can help us find a good oxidizing or reducing agent. All of the substances at the top of the electrochemical series are good oxidizing agents with positive standard reduction potential values, whereas those at the bottom of the electrochemical series are good reducing agents, with negative standard reduction potential values. For example, the F2 electrode with a standard reduction potential of +2.87 volts is a powerful oxidizing agent, whereas the Li+ electrode with a standard reduction potential of -3.05 volts is a strong reducing agent.

2. Electrochemical Cell Standard Emf (E0) Calculation

The standard emf of the cell is the sum of the two half-cell standard reduction potentials:- half-cell reduction and half-cell oxidation-

Eo Cell = Eo red + EoOX

The standard oxidation potential is always represented in terms of reduction potential by convention. As a result, the standard oxidation potential (E∘ox) = - the standard reduction potential ( E∘red). 

Therefore, E∘cell = reduction half cell standard reduction potential) - (standard reduction potential of oxidation half cell)

Because oxidation occurs at the anode and reduction occurs at the cathode. Hence,

Eo cell = Eo cathode - Eo anode

For Example:- 2Ag+ (aq) + Cd → 2Ag + Cd+2 (aq)

The standard reduction potentials are as follows: Ag+/Ag = 0.80 volts, Cd+2/Cd =-0.40 volts.

By the above reaction, we can observe that Cd loses electrons while Ag+ gains. As a result, the oxidation half cell, or anode, is Cd.

Applying the formula,

Eo cell = Eo cathode - Eo anode

                 = 0.80 - (-0.40)

                 = 1.20 volt

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3. Predicting Redox Reaction Possibility

Any redox reaction will occur spontaneously if the free energy change (Go) is negative. Free energy is linked to the cell emf in the following way-

ΔGo = nFEO

Where n represents the number of electrons involved, F represents the Faraday constant, and Eo represents the cell emf.

• If Eo is positive, Go might be negative. 
• When Eo is positive, the cell response is spontaneous and acts as an electrical energy source. 
• If the result is negative, the spontaneous response cannot occur. 
• The calculated Eo value for redox reaction is useful in forecasting the stability of a metal salt solution when kept in another metal container.

Also read more about the Nitride, Spectrochemical Series, and Difference Between Isotopes and Isobars.

4. Predicting the Electrolysis Product

In the presence of two or more types of positive and negative ions in solution, certain metal ions are discharged or freed at the electrodes in preference to others during electrolysis. In such a competition, the ion with the higher oxidizing potential (high standard reduction potential) is discharged first at the cathode. Thus, when an aqueous solution of NaCl containing Na+, Cl-, H+, and OH- ions is electrolyzed, H+ ions are preferentially deposited at the cathode (reduction) rather than Na+ ions, since the reduction potential of hydrogen (0.00 volt) is greater than the reduction potential of sodium (0.00 volt) (-2.71 volts). The anion with the lowest reduction potential will be oxidized at the anode where oxidation occurs. As a result, OH- with a standard reduction potential of 0.40 volts, will be oxidized over Cl- with a standard reduction potential of 1.36 volts.

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Examples of the Electrochemical Series

Example 1 - Fe is lower in the electrochemical series and has a lower reduction potential. It displaces Cu from CuSO₄, whereas Cu is higher in the electrochemical series and has a larger reduction potential. As a result, Fe may easily replace copper in CuSO₄,

Fe + CuSO₄→ FeSO₄ (aq) + Cu i.e.,

Fe + C++(aq) → Fe++(aq) + Cu

Example 2 - Zinc is lower in the electrochemical series and has a lower reduction potential. It displaces Cu from CuSO₄, whereas Cu is higher in the electrochemical series and has a larger reduction potential. As a result, zinc may easily replace copper in CuSO₄,

Zn + CuSO₄→ ZnSO₄ (aq) + Cu i.e.,

Zn + C++(aq) → Zn++(aq) + Cu

Electrochemical Series Important Points

Below are some of the important points to remember for the electrochemical series-

1. The electrochemical series takes an element's reduction potential in relation to the hydrogen scale, where Eo = zero. 
2. The definition of an element's standard reduction potential is "the measure of an element's likelihood to undergo reduction."
3. The larger an element's reduction potential, the easier it is to decrease. 
4. Meanwhile, elements with limited reduction potential oxidize significantly more quickly and readily.
5. Elements with a lower or negative reduction potential, on the other hand, readily give up electrons.
6. Elements having a positive or larger reduction potential quickly accept electrons but do not readily throw them away.
7. In the electrochemical series below hydrogen, stronger reducing agents with a negative standard reduction potential are frequently discovered. 
8. However, weaker reducing agents with positive standard reduction potential are detected above hydrogen in the series.
9. As we travel down the group, the strength of the reducing agent increases while the intensity of the oxidizing agent drops.
10. Similarly, as we progress through the series, the electropositive activity of metals increases. It is reduced in the case of nonmetals.

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Conclusion

The greater an element's reduction potential, the easier it is to decrease. Elements having a low reduction potential, on the other hand, oxidize significantly more quickly and more readily. Electrons are quickly lost in elements having negative or low reduction potentials. Elements having positive or greater reduction potentials do not lose electrons quickly, but they do rapidly gain electrons. In the electrochemical series, stronger reducing agents with a negative standard reduction potential are typically discovered. In this series, however, a weak. Above hydrogen, the reducing agent with a positive standard reduction potential is determined. As you progress through the group, the reducing agent's strength grows while the oxidant's strength diminishes. Similarly, as you progress along the sequence, metal positivity and activity grow or improve. It reduces nonmetals.