Standard Electrode Potential

An example of the standard electrode potential is “the magnitude of the standard emf (electromotive force) of a cell in which molecular hydrogen under standard pressure is oxidised to solvated protons at the left-hand electrode,” according to the definition in electrochemistry. When applied to any element or compound, its reducing power is measured in molar equivalents. It is always a redox reaction that serves as the foundation for an electrochemical cell, such as the galvanic cell, which can be divided into two half-reactions: oxidation at the anode (loss of one electron) and reduction at the cathode (gain of electron). It is the difference in electric potential between the individual potentials of the two metal electrodes with regard to the electrolyte that results in the production of electricity in this situation.

The electric potential fluctuates with temperature, concentration, and pressure, among other factors as well. Due to the fact that the oxidation potential of a half-reaction is the inverse of the reduction potential in a redox reaction, it is sufficient to calculate either of the two potentials in this case. In order to avoid confusion, the standard electrode potential is sometimes referred to as the standard reduction potential. Each electrode-electrolyte interface has a tendency for metal ions from the solution to deposit on the metal electrode as a result of the electrode’s attempt to make it positively charged (deposition). The metal atoms of the electrode have a tendency to dissolve into the solution as ions, leaving the electrons at the electrode behind in an attempt to make the electrode negatively charged. At equilibrium, there is a separation of charges, and depending on the inclinations of the two opposing reactions, the electrode might be either positively or negatively charged in relation to the solution when the reaction is complete. The electrode potential is the difference in potential between the electrode and the electrolyte that arises between the electrode and the electrolyte.

Significance of Standard Electrode Potential

Redox reactions occur in both half-reactions of an electrochemical cell, which is why it is called an electrochemical cell.

It is oxidation that takes place at the anode end, and reduction that takes place at the cathode end. At the anode end, oxidations cause electrons to be lost, whereas at the cathode end, they cause electrons to be gained. As a result, the movement of electrons from anode to cathode results in the conduction of electricity.

Because of the difference in potentials between the cathode and anode of each electrode when it is dipped in its electrolyte, each electrode produces an electric potential. A voltmeter is used to determine the voltage of the cell.

The potentials of individual half-cells cannot be obtained because they are too small. The significance of Standard Electric Potential becomes apparent at this point, because the individual potential can vary depending on changes in pressure, temperature, or concentration of the electrolytes present in the solution. By employing a Standard Hydrogen Electrode, it is possible to determine the individual reduction potential of each half-cell (SHE). When SHE is used, the electrode potential is zero volts.

It is possible to calculate the standard electrode potential of an electrode by attaching an electrode to the SHE and measuring the cell potential of the consequent galvanic cell that is produced. The oxidation potential of an electrode is equal to the inverse of its reduction potential. As a result, the standard electrode potential of an electrode is determined by the standard reduction potential of the electrode.

When it comes to good oxidising agents, their standard reduction potentials are high, whereas when it comes to good reducing agents, their standard reduction potentials are low.

Electrochemical Series

Electrochemical series is the arrangement of elements in accordance with their standard electrode potential values; it is a type of arrangement. It is referred to as an activity series in some circles. The elements with greater standard electrode potentials are put above the elements with lower standard electrode potentials, and the reverse is true. The elements that are placed at the very top of the series have a propensity to be reduced very quickly in size. The elements that are positioned at the bottom, on the other hand, have the least potential to be reduced. Fluorine has the greatest tendency to be reduced due to the fact that it has the highest standard electrode potential of any element. As a result of having the lowest standard electrode potential value, lithium has the least tendency to be decreased of all the elements. As a result, fluorine is an extremely potent oxidising agent, while lithium is an extremely powerful reducing agent.

The spontaneity of Redox Reactions

To be considered spontaneous, the change in Gibbs free energy (ΔGocell) must be negative in the case of a redox reaction. In the following equation, we can see how it works:

ΔG°cell = -nFE°cell

In this equation, n is the total number of moles of electrons required to produce one mole of product, and F denotes Faraday’s constant (approximately 96485 C.mol-1).

The E°cell can be calculated with the assistance of the equation shown below:

E°cell = E°cathode – E°anode

As a result, by subtracting the standard electrode potential of the anode from the standard electrode potential of the cathode, the E°cell can be calculated. To be considered spontaneous, the E°cell must have a positive value (because to the fact that both n and F have positive positive values, and the ΔGocell value must be negative) and the ΔGocell value must be negative.

This indicates that in a spontaneous process,

E°cell is more than zero, which means that E°cathode is greater than E°anode

Consequently, knowing the typical electrode potential of the cathode and anode can aid in forecasting how spontaneously a cell reaction will occur. When it comes to galvanic cells, the  ΔGo of the cell is negative, and when it comes to electrolytic cells, it is positive.

Standard Hydrogen Electrode

Informally abbreviated as SHE, the Standard Hydrogen Electrodes (also known as redox electrodes) serve as the fundamental building block of the oxidation-reduction potential scale. Furthermore, it is the primary electrode at which the electrochemical reaction takes place, hence it is important.

A redox reaction is a chemical reaction in which both oxidation (the gain of an electron by an atom) and reduction (the loss of an electron by an atom) occur at the same time in the same system.

At 25 ° Celsius, the absolute potential of the Standard Hydrogen Electrode is roughly 4.44 0.02 V, which is used in electrochemistry.

When comparing the potential of hydrogen with the potential of other electrochemical reactions, it is said to have a zero volts potential (E°) at a specific temperature (see below) (298 K).

In order to compare the potential of any electrode to the potential of a hydrogen electrode, the temperature of the electrode must be known. It is referred to as the standard hydrogen electrode because it serves as a reference electrode in the electrolysis of hydrogen gas.

Electrochemical cell potential calculation

From the half-reactions and the operating parameters of an electrochemical cell, it is possible to compute the cell potential of the cell.

To obtain the overall standard cell potential, multiply the two half-cell potentials together.

E°cell= E°red + E°ox

Non-standard state conditions

If the conditions are not typical (for example, concentrations greater than 1 mol/L, pressures greater than 1 atm, and temperatures greater than 25°C), we must conduct a few additional procedures.

1. Determine the standard cell potential by measuring the voltage across the cell.
2.  Calculate the new cell potential that results as a result of the altered circumstances.
3.  Calculate the reaction quotient, abbreviated Q.
4. The number of moles electrons exchanged in the reaction is denoted by the letter n in the equation.
5.  Determine Ecell, the cell potential under the non-standard state conditions using the Nernst equation.

The Nernst equation can be written as

Ecell = E°cell – (RT/nF)lnQ.

where

Ecell is the potential of the cell under non-standard state conditions.

E°cell is the potential of the cell in its standard form.

R is the universal gas constant (8.314 JK-1mol-1=8.314 VC-1mol-1) and is equal to 8.314 VC-1mol-1.

T is the Kelvin temperature scale.

In the balanced equation for the cell reaction, n represents the number of moles of electrons transported.

Q is the reaction quotient for the reaction 

Conclusion

It is possible for the same metal to act as an anode in one reaction and as a cathode in another reaction, depending on the electrode potential of each metal. The positive or negative reading on the voltmeter will indicate whether this is the case. The standard cell potential gives us the difference in potential between the cathode and the anode of a cell. When defining electrode potential, the voltage or potential difference between a cell formed from a standard hydrogen electrode and the provided electrode with the potential that is being specified is defined as the electrode potential for that electrode. According to convention, the standard hydrogen electrode is always at a potential of zero volts. The value of the standard electrode potential is zero, and it serves as the foundation for calculating cell potentials when utilising different electrodes or varied quantities of different substances.