What is the Electrochemical Series? How Does it Help in Predicting Whether a Particular Redox Reaction is Feasible in a Given Direction or Not?

The electrochemical series is a fundamental concept in chemistry, serving as a powerful tool to predict the spontaneity and feasibility of redox (reduction-oxidation) reactions. It is essentially a list of chemical species arranged in order of their standard electrode potentials.

Understanding the Electrochemical Series

The electrochemical series, also known as the activity series or electromotive series, is a comprehensive list of elements and their ions organized according to their standard reduction potentials (E°). These potentials are measured under standard conditions (25°C, 1 M concentration for solutions, 1 atm pressure for gases).

Standard Electrode Potential (E°): This value quantifies the tendency of a chemical species to gain electrons (be reduced).
- A more positive E° value indicates a greater tendency for reduction, meaning the species is a stronger oxidizing agent.
- A more negative E° value indicates a greater tendency for oxidation (less tendency for reduction), meaning the species is a stronger reducing agent.
Reference Point: The standard hydrogen electrode (SHE) is used as the universal reference, arbitrarily assigned a standard electrode potential of 0.00 Volts. All other potentials are measured relative to the SHE.

Predicting Redox Reaction Feasibility

The electrochemical series provides a clear framework for predicting whether a specific redox reaction will proceed spontaneously in a given direction. This prediction is based on the calculation of the standard cell potential (E°cell) for the proposed reaction.

The Key Principle:

A redox reaction is considered feasible (or spontaneous) under standard conditions if the calculated standard cell potential (E°cell) for that reaction is positive (> 0).

How it Works:

Identify Half-Reactions: Break down the overall redox reaction into its two constituent half-reactions: one reduction and one oxidation.
Assign Standard Reduction Potentials: Look up the standard reduction potentials (E°) for both half-reactions from the electrochemical series.
Determine Cathode and Anode:
- The species with the higher (more positive) standard reduction potential will be reduced (gain electrons) and act as the oxidizing agent. This half-reaction occurs at the cathode.
- The species with the lower (more negative) standard reduction potential will be oxidized (lose electrons) and act as the reducing agent. This half-reaction occurs at the anode.
- A substance with a higher standard reduction potential possesses a greater tendency to gain electrons and thus can oxidize another substance.
Calculate Standard Cell Potential (E°cell):
The standard cell potential is calculated using the formula:

E°cell = E°(reduction at cathode) - E°(reduction at anode)

Alternatively, it can be expressed as:

E°cell = E°(reduction half-reaction) + E°(oxidation half-reaction, where E° for oxidation is the negative of its standard reduction potential)
Interpret the Result:
- If E°cell > 0: The reaction is feasible and will proceed spontaneously in the direction written. This indicates that the Gibbs free energy change (ΔG°) for the reaction is negative (ΔG° = -nFE°cell, where n is moles of electrons, F is Faraday constant).
- If E°cell < 0: The reaction is not feasible in the direction written. Instead, the reverse reaction would be spontaneous.
- If E°cell = 0: The reaction is at equilibrium.

Practical Examples:

Let's illustrate with common metal displacement reactions:

Example 1: Will Zinc react with Copper Sulfate?
Consider the reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Half-reactions and E° values:
- Reduction: Cu²⁺(aq) + 2e⁻ → Cu(s) ; E° = +0.34 V
- Oxidation: Zn(s) → Zn²⁺(aq) + 2e⁻ ; E° = -0.76 V (standard reduction potential for Zn²⁺ + 2e⁻ → Zn is -0.76 V, so for oxidation it's +0.76 V)
Determine cathode/anode: Copper has a higher reduction potential (+0.34 V) than Zinc (-0.76 V). Therefore, Cu²⁺ will be reduced (cathode), and Zn will be oxidized (anode).
Calculate E°cell:
E°cell = E°(Cu²⁺/Cu) - E°(Zn²⁺/Zn)
E°cell = (+0.34 V) - (-0.76 V) = +1.10 V
Feasibility: Since E°cell is positive (+1.10 V), the reaction Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) is feasible and will proceed spontaneously. This is why zinc metal will displace copper from a copper sulfate solution.

Example 2: Will Copper react with Zinc Sulfate?
Consider the reverse reaction: Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s)

Half-reactions and E° values:
- Reduction: Zn²⁺(aq) + 2e⁻ → Zn(s) ; E° = -0.76 V
- Oxidation: Cu(s) → Cu²⁺(aq) + 2e⁻ ; E° = +0.34 V (standard reduction potential for Cu²⁺ + 2e⁻ → Cu is +0.34 V, so for oxidation it's -0.34 V)
Determine cathode/anode: Zinc has a lower reduction potential (-0.76 V) compared to Copper (+0.34 V). For the reaction to proceed as written, Zn²⁺ would need to be reduced, and Cu oxidized.
Calculate E°cell:
E°cell = E°(Zn²⁺/Zn) - E°(Cu²⁺/Cu)
E°cell = (-0.76 V) - (+0.34 V) = -1.10 V
Feasibility: Since E°cell is negative (-1.10 V), the reaction Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s) is not feasible in this direction. Copper metal will not displace zinc from a zinc sulfate solution.

Other Important Applications:

Designing Electrochemical Cells/Batteries: The electrochemical series helps in selecting appropriate electrode materials to generate a desired voltage.
Corrosion Prevention: Understanding electrode potentials helps predict which metals are more susceptible to corrosion and guides the use of sacrificial anodes or protective coatings.
Electroplating: It's crucial for determining which metal will deposit onto a surface under specific conditions.
Predicting Reactivity: Metals higher in the series (more negative E°) are more reactive as reducing agents, while non-metals with higher E° are more reactive as oxidizing agents.

In conclusion, the electrochemical series provides a robust quantitative method for predicting the feasibility and direction of redox reactions, making it an indispensable tool in chemistry and related fields.