Standard Electrode Potential (E°)
What is Electrode Potential?
Imagine two different metal strips, like zinc and copper, placed in their own salt solutions. If you connect them with a wire, a tiny electric current flows. Why? Because each metal has a different inherent "desire" to lose electrons and become a positive ion. This inherent tendency is the electrode potential.
When a metal rod (an electrode) is dipped into a solution containing its own ions (e.g., a Zn rod in ZnSO₄ solution), one of two things happens:
1. The metal atoms may lose electrons and enter the solution as positive ions: $ Zn_{(s)} \rightarrow Zn^{2+}_{(aq)} + 2e^- $. This leaves the metal rod with a negative charge.
2. The positive ions from the solution may gain electrons and get deposited on the metal rod: $ Cu^{2+}_{(aq)} + 2e^- \rightarrow Cu_{(s)} $. This leaves the metal rod with a positive charge.
In both cases, a potential difference is created between the metal electrode and its solution. This is the electrode potential. However, we cannot measure the absolute potential of a single half-cell; we can only measure the difference in potential between two half-cells. This is why we need a universal reference point.
The Standard Hydrogen Electrode: The Universal Reference
To create a consistent scale for comparing all electrode potentials, scientists established the Standard Hydrogen Electrode (SHE)[1] as the baseline. By definition, the standard electrode potential of the SHE is exactly 0.00 V.
To find the Standard Electrode Potential (E°) of any other half-cell, it is connected to the SHE. The voltage measured across the two electrodes is the E° value for that half-cell. The conditions must be standard: a temperature of 298 K (25 °C), a pressure of 100 kPa for any gases involved, and a solution concentration of exactly 1 mol dm⁻³.
The Standard Electrode Potential Series
When we measure the E° values for various half-cells and list them in order, we get the Electrochemical Series. This series is a powerful tool for predicting the spontaneity of redox reactions.
| Half-Cell Reaction | E° (Volts) |
|---|---|
| $ Li^+ + e^- \rightarrow Li_{(s)} $ | -3.04 |
| $ Zn^{2+} + 2e^- \rightarrow Zn_{(s)} $ | -0.76 |
| $ 2H^+ + 2e^- \rightarrow H_{2(g)} $ (SHE) | 0.00 |
| $ Cu^{2+} + 2e^- \rightarrow Cu_{(s)} $ | +0.34 |
| $ Ag^+ + e^- \rightarrow Ag_{(s)} $ | +0.80 |
| $ F_{2(g)} + 2e^- \rightarrow 2F^- $ | +2.87 |
Key Interpretations of the Series:
• Negative E° Values: Elements like Lithium (Li) and Zinc (Zn) have highly negative E° values. This means they have a strong tendency to lose electrons and undergo oxidation. They are strong reducing agents.
• Positive E° Values: Elements like Silver (Ag) and especially Fluorine (F₂) have highly positive E° values. This means they have a strong tendency to gain electrons and undergo reduction. They are strong oxidizing agents.
• The "Noble" Metals: Metals with positive E° values, such as silver, gold, and platinum, are found at the bottom of the series. They are unreactive and do not corrode easily, which is why they are called "noble" metals.
Predicting Cell Voltage and Reaction Spontaneity
The primary application of E° values is to predict the voltage of a galvanic (voltaic) cell and whether a redox reaction will occur spontaneously.
Rule 1: The Spontaneity Rule
A redox reaction is spontaneous (will happen on its own) if the E° value for the reduction half-reaction is more positive than the E° value for the oxidation half-reaction. In simpler terms, the half-cell with the higher (more positive) E° value will undergo reduction, and the one with the lower (more negative) E° value will undergo oxidation.
Rule 2: Calculating Cell Potential (E°cell)
The standard electromotive force (EMF) of a cell is calculated as:
Where the cathode is the electrode where reduction occurs, and the anode is the electrode where oxidation occurs. A positive E°cell indicates a spontaneous reaction.
A Practical Example: The Lemon Battery
Let's build a classic lemon battery with a zinc nail and a copper coin. According to the electrochemical series:
• Zinc half-reaction: $ Zn^{2+} + 2e^- \rightarrow Zn_{(s)} $ E° = -0.76 V
• Copper half-reaction: $ Cu^{2+} + 2e^- \rightarrow Cu_{(s)} $ E° = +0.34 V
Since copper has the more positive E° value, it will act as the cathode (reduction will occur here: $ Cu^{2+} + 2e^- \rightarrow Cu $). Zinc, with the more negative E° value, will act as the anode (oxidation will occur here: $ Zn \rightarrow Zn^{2+} + 2e^- $).
The overall cell reaction is: $ Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Cu_{(s)} $
We can calculate the theoretical voltage:
$ E°_{cell} = E°_{cathode} - E°_{anode} = E°_{Cu} - E°_{Zn} = 0.34 - (-0.76) = +1.10 V $
This positive value confirms that the reaction is spontaneous, and our lemon battery will produce a voltage close to 1.1 V, enough to power a small LED or a digital clock.
Important Questions
Voltage is a measure of potential difference, just like height is a measure relative to sea level. You can't have a "height" without a reference point. Similarly, we need a reference electrode (the SHE) to measure the "electron pressure" or potential of any other half-cell against it.
Generally, yes. Metals with negative E° values (like Magnesium, Aluminum, Zinc) are more susceptible to corrosion (oxidation) because they have a high tendency to lose electrons. This is why zinc is used to galvanize (coat) steel; the zinc corrodes first, sacrificially protecting the iron.
A reaction with a negative E°cell is non-spontaneous under standard conditions. However, it can be forced to occur by supplying electrical energy from an external source. This is the principle behind electrolysis, used for processes like electroplating and recharging batteries.
The concept of Standard Electrode Potential (E°) provides a quantitative and predictive framework for the entire field of electrochemistry. By serving as a universal measuring stick for the tendency of species to gain or lose electrons, it allows us to understand why batteries work, predict which metals will corrode, design industrial electrolysis processes, and grasp the fundamental energy changes in redox reactions. From the simple lemon battery to the complex lithium-ion battery in a smartphone, the principles governed by the electrochemical series are silently at work, powering our modern world.
Footnote
[1] SHE (Standard Hydrogen Electrode): The universal reference electrode with a defined potential of 0.00 V, against which all other standard electrode potentials are measured.
[2] EMF (Electromotive Force): The maximum potential difference between the electrodes of a galvanic cell when no current is flowing. It is a measure of the cell's voltage.
[3] Redox Reaction: A chemical reaction involving the simultaneous processes of reduction (gain of electrons) and oxidation (loss of electrons).