Standard Electrode Potential

Introduction

The standard electrode potential, denoted as E°, is a measure of the intrinsic tendency of a redox system to lose or gain electrons. It is a thermodynamic quantity that indicates how easily a substance can be oxidized or reduced. The standard electrode potentials are measured under standard conditions, which are:

• Temperature of 25°C (298 K)
• Pressure of 1 atmosphere (for gases)
• Concentration of 1 M (for solutions)

These potentials are measured relative to the standard hydrogen electrode (SHE), which is assigned a potential of 0.00 volts.

Understanding Standard Electrode Potential

The standard electrode potential is determined by the Gibbs free energy change (ΔG°) for the redox reaction. The relationship between ΔG° and E° is given by the equation:

$$ \Delta G° = -nFE° $$

where:

• ΔG° is the standard Gibbs free energy change in joules (J)
• n is the number of moles of electrons transferred
• F is the Faraday constant (approximately 96485 C/mol)
• E° is the standard electrode potential in volts (V)

A positive E° value indicates a tendency to gain electrons and be reduced (a strong oxidizing agent), while a negative E° value indicates a tendency to lose electrons and be oxidized (a strong reducing agent).

Standard Hydrogen Electrode (SHE)

The standard hydrogen electrode consists of a platinum electrode coated with platinum black, in contact with 1 M H⁺ ions and bathed by hydrogen gas at 1 atm. The half-reaction is:

$$ 2H⁺(aq) + 2e^- \leftrightarrow H_2(g) $$

The SHE serves as the reference electrode against which all other electrode potentials are measured.

Table of Standard Electrode Potentials

Here is a table of some common standard electrode potentials:

Half-Reaction	Standard Electrode Potential (E°)
F₂(g) + 2e⁻ → 2F⁻(aq)	+2.87 V
O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)	+1.23 V
Cl₂(g) + 2e⁻ → 2Cl⁻(aq)	+1.36 V
2H⁺(aq) + 2e⁻ → H₂(g)	0.00 V
Zn²⁺(aq) + 2e⁻ → Zn(s)	-0.76 V
Fe²⁺(aq) + 2e⁻ → Fe(s)	-0.44 V
Cu²⁺(aq) + 2e⁻ → Cu(s)	+0.34 V
Ag⁺(aq) + e⁻ → Ag(s)	+0.80 V

Calculating Cell Potential

The standard cell potential (E°_cell) for a galvanic cell can be calculated using the standard electrode potentials of the two half-cells:

$$ E°{cell} = E°{cathode} - E°_{anode} $$

The cathode is where reduction occurs, and the anode is where oxidation occurs.

Example

Consider a galvanic cell with a zinc electrode in a Zn²⁺ solution and a copper electrode in a Cu²⁺ solution. The standard electrode potentials are:

• Zn²⁺(aq) + 2e⁻ → Zn(s) : E° = -0.76 V (anode)
• Cu²⁺(aq) + 2e⁻ → Cu(s) : E° = +0.34 V (cathode)

The standard cell potential is:

$$ E°{cell} = E°{cathode} - E°_{anode} = +0.34 V - (-0.76 V) = +1.10 V $$

Nernst Equation

The Nernst equation allows us to calculate the electrode potential under non-standard conditions:

$$ E = E° - \frac{RT}{nF} \ln Q $$

where:

• E is the electrode potential under non-standard conditions
• R is the universal gas constant (8.314 J/(mol·K))
• T is the temperature in Kelvin
• n is the number of moles of electrons transferred
• F is the Faraday constant
• Q is the reaction quotient

For reactions at 25°C (298 K), the Nernst equation simplifies to:

$$ E = E° - \frac{0.0592}{n} \log Q $$

Example

For the half-reaction Cu²⁺(aq) + 2e⁻ → Cu(s) at 25°C, if the concentration of Cu²⁺ is 0.01 M, the electrode potential can be calculated as:

$$ E = E° - \frac{0.0592}{n} \log \frac{1}{[Cu^{2+}]} $$ $$ E = 0.34 V - \frac{0.0592}{2} \log \frac{1}{0.01} $$ $$ E = 0.34 V - 0.0296 \log 100 $$ $$ E = 0.34 V - 0.0296 \times 2 $$ $$ E = 0.34 V - 0.0592 V $$ $$ E = 0.2808 V $$

Conclusion

The standard electrode potential is a fundamental concept in electrochemistry that helps predict the direction of redox reactions and the feasibility of electrochemical cells. It is essential for understanding the behavior of batteries, electrolysis, and corrosion processes. By using the Nernst equation, we can also predict the behavior of electrochemical systems under non-standard conditions.