Reduction Potential

Reduction potential measures a chemical species’ tendency to gain or lose electrons from an electrode and thus reduce or oxidise. We can express the redox potential in volts or millivolts.

It means that every species contains unique inherent redox potential. For example, a species with a greater reduction potential has a greater affinity for electrons and is thus more likely to reduce. 

Redox potential measures the environmental conditions’ effect on the chemical or electrochemical reactivity of substances. Therefore, we can also use it for predicting the corrosion protection of various substances and systems.

Following are the reduction potential examples:

Measurement and Interpretation

Redox potential tends to gain or lose electrons when introducing a new species into an aqueous solution. Therefore, positive reduction potential solutions will accept new species of electrons more easily. We can define reduction potentials with a reference electrode because measuring the absolute potentials is practically impossible. Instead, we can measure the potential difference between a stable reference electrode connected by a salt bridge and an inert sensing electrode to determine the aqueous solution’s reduction potential.

We can use an electron transfer platform made of platinum, gold, or graphite to transfer electrons to or from the reference half-cell. The reference half-cell has a known redox standard potential. The standard hydrogen electrode (SHE), which assigns an arbitrary half-cell potential of 0.0 mV, serves as a reference for all standard redox potentials. But it’s fragile and impractical for everyday lab use. We use silver chloride (SC) and saturated calomel (SCE) as more stable reference electrodes.

Although measuring the aqueous solution’s redox potential is very straightforward. Several issues limit the aqueous solution’s redox potential interpretation, including non-equilibrium, several redox couples presence, poisoning of the electrode, currents small exchange, solution temperature’s effects and pH, slow electrode kinetics, irreversible reactions, and inert redox couples. As a result, practical measurements rarely match calculated values. However, the reduction potential measurement has proven beneficial as an analytical tool for monitoring system changes rather than determining their absolute value.

Reduction Potential Explanation

A potential difference forms at the metal-solution interface when a metal immerses in its ion solution. The magnitude of the potential difference reveals the electrodes’ tendency to oxidise, lose, or acquire electrons.

The ion and metal represent half-cells, and a reaction is considered a half-reaction. Immersed metal is called an electrode. The potential after the reaction of the solution and electrode is known as the electrode potential. Here, the losing or gaining of electrons is known as electrode potential. Reduction happening at the electrode is known as reduction potential.

The oxidation potential occurs when oxidation occurs at the electrode:

M ⇢ M₂+ + 2e^–

A positive charge builds up on the metal rod as metal ions begin to deposit on the metal surface. Since oxidation is the inverse of reduction, changing the sign provides the reduction potential.

For an electrode, in general,

Oxidation potential = –Reduction potential

As an example, the standard oxidation potential for zinc is

Eo (Zn/Zn²⁺) = 0.76V

as well as standard reduction potential as 

Eo (Zn²⁺/Zn) = -0.76V

All electrode potentials are usually expressed as reduction potentials.

The IUPAC (International Union of Pure and Applied Chemistry) recently adopted the reduction potential as a designation for electrode potential.

We can represent the standard electrode potential by E^o when we carry out the half cell reaction at 298 K, and the electrode suspends in a one molar concentration solution. Using the standard electrode potential E^o, we may evaluate the thermodynamic activity of various chemical substances. But there’s no method to determine its absolute value. We can measure an electrode’s potential using a hydrogen electrode.

The concentration in solution in contact with metal determines the electrode potential. When the concentration of ions increases, the oxidation potential of an electrode decreases, and when the concentration of ions decreases, the reduction potential increases.

Half Cells

An electrolyte separates the two half cells in a battery, making the battery a cell. The electrodes connect the half cells to a circuit outside the cell. Although it isn’t necessary, each electrode can be part of a redox couple.

H+(aq) concentration, hydrogen gas pressure (105Pa), and temperature (298) K are the standard operating parameters for the hydrogen half cell.

We may measure each half cell to the reference hydrogen half cell, creating a list of electrode potentials for a standard hydrogen half cell. This half cell half-reaction is

2H⁺(aq) + 2e^– ⇌ H₂(g)

Because the potentials of the electrodes vary with temperature, we define a standard temperature, and that’s 298K. Because changing the concentration of any ions in the half-reactions affects the voltages, we use a standard ion concentration of 1.00 mol dm^-3. The standard pressure is 105 Pa.

• A standard hydrogen half-cell has a potential of 0.0V, chosen for convenience.
• The difference between half cell and standard hydrogen half cell is the standard electrode potential E^o.
• The Eo Values change sign when dipped in a 1.00 mol dm^-3 metal salt solution at 298K.

Conclusion

The ability to gain electrons and lose them is the reduction potential. ‘Volts or millivolts are the units used to express this voltage. Positive reduction potential values indicate a greater tendency to reduce. The metal’s ionic activity decreases as a result of the complex formation. Thus, complexation reduces the metal ion’s reduction potential.

The standard reduction potential for a free Co⁺³ ion is 1.853 V; however, in the complexed state [Co(NH₃)₆]⁺³, the standard reduction potential lowers to 0.1 V. [Fe(CN)₆]^-3 reduction potential is 0.36 V. In comparison, the standard reduction potential of free Fe⁺³ is 0.771 V.

In general, complex formation reduces the metal ion’s reduction potential, which indicates that higher oxidation states of the metal ion are more stable after complexation. Therefore, we can use this change in reduction potential to detect complex formation in synthetic coordination chemistry.